Surface Complexes Of Lead And Organic Acids At The Hematite / Water Interface
DISSERTATION
Presented in Partial Fulfillment of the Requirements for the Degree Doctor of Philosophy in the Graduate School of The Ohio State University
By
Matthew Noerpel
Graduate Program in Civil Engineering
The Ohio State University
2015
Dissertation Committee:
Professor John J. Lenhart, Advisor
Professor Heather Allen
Professor Yu-Ping Chin
Professor Linda Weavers ! ! ! ! ! ! ! ! ! ! ! ! ! ! ! ! Copyright by
Matthew Noerpel
2015
! ! ! ! ! !
Abstract
Lead is a common and very toxic contaminant in the environment. Consumption of lead by children can cause irreversible harm to the brain and central nervous system. It is crucial to understand the behavior of lead in the environment in order to protect the population from coming to harm. Colloidal iron oxide particles and organic acids are ubiquitous in the natural environment. In combination and independently, they play an important role in the fate of lead and other heavy metal contaminants. Lead can adsorb onto the surface of the particles and remain mobile as the small particles do not settle out of suspension. Organic acids can adsorb on the surface of the mineral particles changing their surface charge, stability and reactivity as well as interacting with lead in solution. It is therefore important to understand the interactions of organic acids and colloidal particles with and without lead in order to fully understand the fate of lead in the environment.
Throughout this project, the iron oxide hematite was used as the adsorbent mineral phase.
In the first chapter, we investigated the adsorption mechanisms that bond the common organic acid, citric acid, to the hematite surface using batch adsorption, Fourier transform infrared spectroscopy (FTIR), and molecular modeling and surface complexation modeling (SCM). All of the methods used indicated that the dominant adsorption mode is as an outer-sphere complex that changes protonation state with pH, going from singly
ii protonated at low pH to deprotonated at higher pH conditions. There was also evidence of an inner-sphere bidentate complex at low pH.
In Chapter 3, the adsorption of lead on bare hematite particles and single crystal surfaces was examined using two synchrotron based X-ray techniques, extended X-ray adsorption fine structure (EXAFS), on particles, and X-ray reflectivity (XR) on single crystal surface with a known surface exposed. The results of the two techniques confirm that lead adsorbs as an inner-sphere bidentate complex in an edge-sharing and corner-sharing arrangement. In addition, the XR method found an additional outer-sphere complex on formed on the single crystal surfaces that is not visible using EXAFS. The reactivity of the three surfaces tested varied greatly. The surface that is considered the most common face, the (001), was the least reactive face and the adsorption did not change with pH.
In Chapter 4, the same X-ray techniques were used to determine the influence that organic acids have on the surface. Four acids were tested, citrate, phthalate, humic and fulvic acids. The results showed that at pH 4, the acids enhanced the adsorption of lead onto the particle surfaces, however at pH 6, citrate hindered the adsorption of lead as it likely chelated the lead and held it in solution. The single crystal studies showed that in the presence of organic acids, the lead still adsorbs strongly as an inner-sphere complex.
Overall, the specific surface played a larger role in determining the manner in which the lead adsorbed than did the acid added.
iii
Dedication
To Dana and Josephine
iv
Acknowledgments
First I would like to thank my advisor, Dr. John Lenhart. His guidance, encouragement and patience with matters both in and outside of the lab have been invaluable. Dr.
Lenhart allowed me the freedom to explore new methods that made the last five years a more challenging and exciting period than I could ever have imagined before coming to
Ohio State.
I would like to thank my committee members, Drs. Heather Allen, Paula Mouser, Yu- ping Chin, and Linda Weavers for showing interest in my work and challenging me.
Thanks also to Jason Cheng and Zuzana Bohrerova for their help in the lab.
The third and fourth chapters of this dissertation would not have been possible if it were not for the aid of Dr. Sang Soo Lee. His knowledge and excitement at the synchrotron made the marathon data collection sessions possible, and his patience in explaining both the theory of the reflectivity method and how to deal with the many gigabytes of data we collected were critical to the success of the single crystal experiments.
I also owe a debt of gratitude to all of the other grad students in the lab for broadening my knowledge with the many discussions of a wide variety of topics from sonication processes to the probable presence and importance of algae in the clouds. Our many discussions added insight and levity to the long days in the basement. Thanks to Zongsu,
v Ray, Mike, Mengling, Xuan, Ryan, Katie, Marcia, Jenny, Mary, Yen-Ling, Chenyi, Dan,
Wei and all the others. Thanks also to the undergraduate students I had the pleasure of working with, Drew, Craig, Stephanie, Stephanie, and Jake.
Finally, I would like to thank my family for the constant and unwavering support, interest, and encouragement. When instilling a love of education in me, my parents probably didn’t expect me to stay in school until I was 33. Thanks to my Kuhnline in- laws, for providing me with all the computer equipment I needed to write this dissertation. And the biggest thanks of course to my wife Dana and daughter Josephine for all the patience, joy, and inspiration they provided.
vi
Vita
2004...... B.S. Civil Engineering, Virginia Polytechnic
Institute and State University
2010...... M.S. Environometnal Engineering, Carnegie
Mellon University
2010 to present ...... Graduate Research Associate, Department
of Civil, Environmental and Geodetic
Engineering, The Ohio State University
Publications
Xiao,&R.,&Noerpel,&M.,&Ling&Luk,&H.,&Wei,&Z.,&Spinney,&R.,&2013.&Thermodynamic&and& kinetic&study&of&ibuprofen&with&hydroxyl&radical:&A&density&functional&theory& approach.&International&Journal&of&Quantum&Chemistry&114,&74P83.&
Noerpel, M. and Lenhart, J., 2015. The Impact of Particle Size on the Adsorption of Citrate to Hematite. Journal of Colloid and Interface Science. In Press.
Fields of Study
Major Field: Civil Engineering
vii
Table of Contents
Abstract ...... ii!
Acknowledgments ...... v!
Vita ...... vii!
Table of Contents ...... viii!
List of Tables ...... xiii!
List of Figures ...... xvi!
Chapter 1: Introduction ...... 1!
1.1 Problem Description ...... 1!
1.1.1 Organic acid adsorption ...... 3!
1.1.2 Lead adsorption ...... 5!
1.1.3 Interaction of metals and organic acids ...... 8!
1.2 Research objectives ...... 9!
1.3 Dissertation overview ...... 10!
1.3.1 The impact of particle size on the adsorption of citrate to hematite ...... 10!
1.3.2 X-ray analysis of lead adsorbed on the hematite (001), (012), and (110) surface
...... 10! viii 1.3.3 Lead and organic acids on hematite ...... 11!
References ...... 15!
Chapter 2: The Impact of Particle Size on the Adsorption of Citrate to Hematite ...... 23!
Abstract ...... 23!
2.1 Introduction ...... 24!
2.2 Materials and Methods ...... 29!
2.2.1 Hematite Synthesis and Characterization ...... 29!
2.2.2 Batch Adsorption Protocol ...... 31!
2.2.3 ATR-FTIR Spectroscopy ...... 32!
2.2.4 Molecular Modeling ...... 33!
2.2.5 Surface Complexation Modeling ...... 34!
2.3 Results and Discussion ...... 35!
2.3.1 Adsorption data ...... 35!
2.3.2 FTIR spectroscopy ...... 37!
2.3.3 Computational Modeling ...... 42!
2.3.4 Surface complexation modeling ...... 44!
2.4 Conclusion ...... 49!
References ...... 51!
ix Chapter 3: X-Ray Analysis Of Lead Adsorbed On The Hematite (001), (012), And (110)
Surface ...... 68!
3.1 Introduction ...... 69!
3.2 Experimental ...... 72!
3.2.1 Hematite ...... 72!
3.2.2 EXAFS ...... 73!
3.2.3 X-ray Reflectivity ...... 75!
3.3 Results and Discussion ...... 77!
3.3.1 EXAFS ...... 77!
3.3.2 (001) Surface ...... 79!
3.3.3 (012) Surface ...... 82!
3.3.4 (110) Surface ...... 86!
3.4 Conclusion ...... 88!
References ...... 90!
Chapter 4: Effect of Organic Acids on Lead Adsorption on Hematite ...... 101!
Abstract ...... 101!
4.1 Introduction ...... 102!
4.2 Methods and materials ...... 109!
4.2.1 Particle synthesis ...... 109!
x 4.2.2 Batch Adsorption ...... 110!
4.2.3 EXAFS ...... 111!
4.2.4 Single Crystal Reflectivity ...... 113!
4.3 Results ...... 115!
4.3.1 EXAFS and batch adsorption ...... 115!
4.3.2 Citrate ...... 116!
4.3.3 Phthalate ...... 120!
4.3.4 Fulvic Acid ...... 124!
4.3.5 Humic acid ...... 127!
4.3.6 General Discussion ...... 129!
References ...... 137!
Chapter 5: Conclusions and Future Work ...... 164!
5.1 Objectives ...... 164!
5.2 Future Work ...... 167!
Appendix A: Supporting Information for Chapter 2 ...... 170!
Appendix B: Crystal description and solution speciation diagrams ...... 185!
Appendix C: X-ray Reflectivity ...... 190!
C.2 Data Collection ...... 191!
C.3 XR Modeling ...... 192!
xi C.4 RAXR Modeling ...... 195!
Appendix D: Comparison of (001), (012), and (110) RAXR ...... 209!
Bibliography ...... 213!
xii
List of Tables
Table 1.1. Summary of recent / relevant studies on the adsorption of citrate on mineral
surfaces...... 12!
Table 1.2. Summary of single crystal hematite experiments...... 13!
Table 1.3. Summary of recent studies on the adsorption of lead to mineral surfaces in the
presence of organic aids or other anions...... 14!
Table 2.1. Constants used for triple layer model fitting. All values taken from Hwang
and Lenhart(Hwang and Lenhart, 2008) aside from the results of the best model
fit of 1000 µM Citrate to the LSA hematite (eq 5-7). Values given as logs of
equilibrium constants...... 65!
Table 2.2. DFT assignments of simulated peaks for citrate and citric acid ...... 66!
Table 2.3. Experimental and theoretical symmetric and asymmetric stretch peak
locations. The protonated mononuclear tridentate structure did not exhibit either a
symmetric or asymmetric stretch. Structures for theoretical complexes are show
in Figure A.2...... 67!
Table 3.1. Fits to Pb LIII EXAFS...... 99!
Table 3.2. Best fit model parameters from the RAXR data of lead only on the three
hematite surfaces...... 100!
Table 4.1. Results of the EXAFS model fitting ...... 158!
xiii Table 4.2. Result from the linear combination fitting ...... 159!
Table 4.3. Results of Model dependent RAXR fit of lead on three faces of hematite in
the presence of citric acid...... 160!
Table 4.4. Results of Model dependent RAXR fit of lead on three faces of hematite in
the presence of phthalic acid...... 161!
Table 4.5. Results of Model dependent RAXR fit of lead on three faces of hematite in
the presence of fulvic acid. At pH 4, the (012) surface required an additional peak
to fit the data...... 162!
Table 4.6. Results of Model dependent RAXR fit of lead on three faces of hematite in
the presence of Humic acid...... 163!
Table A.1. Solution species chemistry used for the surface complexation modeling. .. 175!
Table A.2. Coordinates for protonated outer-sphere complex shown in figure A.2. A . 176!
Table A.3. Coordinates for deprotonated outer-sphere complex shown in figure A.2 B.
...... 177!
Table A.4. Coordinates for protonated inner-sphere mononuclear bidentate complex
(CT-MN) shown in figure A.2 C ...... 178!
Table A.5. Coordinates for protonated inner-sphere mononuclear bidentate complex
(CT-MN) shown in figure A2 D...... 179!
Table A.6. Coordinates for protonated inner-sphere binuclear bidentate complex with the
terminal and central carboxyl groups bound to the iron oxide cluster (CT-BN)
shown in figure A.2 E...... 180!
xiv Table A.7. Coordinates for deprotonated inner-sphere binuclear bidentate complex with
the terminal and central carboxyl groups bound to the iron oxide cluster (CT-BN)
shown in figure A.2 F...... 181!
Table A.8. Coordinates for protonated inner-sphere binuclear bidentate complex with the
central carboxyl group and deprotonated hydroxyl bound to the iron oxide cluster
(CH-BN) shown in figure A.2 G...... 182!
Table A.9. Coordinates for deprotonated inner-sphere binuclear bidentate complex with
the central carboxyl group and deprotonated hydroxyl bound to the iron oxide
cluster (CH-BN) shown in figure A.2 H...... 183!
Table C.1. Range of q over which data was collected for the XR experiments on each
surface ...... 197!
Table C.2. Physical parameters of the three surfaces. The d spacing was found to vary
slightly between beamtimes and is given here to the maximum precision all the
measured values agree to. When the lattice height is off even to the fourth
decimal of precision, the Bragg peak location will change and can be seen in data
points near the Bragg peak...... 198!
Table C.3. Input parameter file for (012) model...... 199!
Table C.4. Input parameter file for (001) model ...... 200!
Table C.5. Input parameter file for (110) model...... 201!
xv
List of Figures
Figure 2.1. Adsorption envelope for citrate (1000 µM) adsorbed onto LSA and HSA
hematite at saturated conditions presented as a) fraction adsorbed and b) surface
coverage...... 58!
Figure 2.2. Surface coverage of adsorbed citrate on the low surface area hematite as a
function of pH at 3 different NaCl concentrations...... 59!
Figure 2.3. Reference spectra used for comparison to adsorbed citrate spectra. The ferric
citrate spectra at 3 pH values represent models for inner-sphere complexation.
Additional experimental spectra (top, black) are presented for fully protonated
citric acid (pH 2.5) and fully deprotonated citrate (pH 9.5). Corresponding
theoretical spectra are also presented (blue, bottom) where the vertical lines (red)
represent the individual absorbance frequencies from the DFT calculation...... 60!
Figure 2.4. Spectra for adsorbed citrate on LSA and HSA hematite at given pH values (1
mM citrate)...... 61!
Figure 2.5. FTIR spectra of citrate adsorbed on LSA hematite at (a) pH 3.0 and (b) 6.5.
The citrate concentrations for the spectra from bottom to top for both pH values
were 62.5, 125, 250, 500 µM...... 62!
Figure 2.6. Selected theoretical infrared spectra for adsorbed citrate structures visualized
with a Lorentzian distribution with a 20 cm-1 FWHM. The short red vertical lines
xvi represent the absorption energies. The dashed vertical lines reflect the peak
locations for the experimental results of the low surface area hematite at mildly
acidic pH. (OS = outersphere, MN=mononuclear, BN = Binuclear, DP =
deprotonated, CT=central and terminal carboxyl bonding, CH=Central and
hydroxyl bonding.) The optimized structures are shown in supporting information
(Figure A.2) ...... 63!
Figure 2.7. Results of surface complexation modeling using one inner-sphere (IS), one
singly protonated outer-sphere (OSP) and one deprotonated outer-sphere (OSDP)
complex as described in eq 5-7 of table 2.1. Individual species are shown with
dashed lines and the total adsorption with a solid line. Equilibrium constants were
determined from the data in plot A and applied to data collected at other solution
conditions (B-D)...... 64!
Figure 3.1. Result of the EXAFS experiments. A) k3 weighted χ(k) functions and B)
their Fourier transform. The data is shown with the blue circles and the fit with
the red line...... 95!
Figure 3.2. Results of the XR / RAXR experiments for lead on the (001) surface of
hematite at pH 4 and pH 6 as noted in the plots. The solid black line is the overall
electron density (XR) and the red area is the lead specific electron density
(RAXR). The blue dashed line is the electron density of the (001) surface in DI
water...... 96!
Figure 3.3. Results of the XR / RAXR experiments for lead on the (012) surface of
hematite at pH 4 and pH 6 as noted in the plots. The solid black line is the overall
electron density (XR) and the read area is the lead specific electron density xvii (RAXR). The blue dashed line is the electron density of the (001) surface in DI
water. The tall XR peak above a height of 0 is the half layer termination...... 97!
Figure 3.4. Results of the XR / RAXR experiments for Lead on the (110) surface of
hematite at pH 4 and pH 6 as noted in the plot. The solid black line is the overall
electron density (XR) and the read area is the lead specific electron density
(RAXR). The blue dashed line is the electron density of the (001) surface in DI
water...... 98!
Figure 4.1. Surface coverage of lead on LSA and HSA hematite particles in the presence
and absence of citric acid...... 145!
Figure 4.2. Lead adsorption results for the EXAFS samples on the HSA hematite ...... 146!
Figure 4.3. EXAFS of standards used for the LCF. The vertical line at k = 3.95 Å-1 is the
antinode location of the lead only adsorbed on hematite and the vertical line at k =
3.45 Å-1 is the location of the Pb-Citrate antinode. Data in blue, fit in red...... 147!
Figure 4.4. EXAFS of citrate and lead adsorbed on LSA and HSA hematite as a function
of pH. Data in blue, fit in red. The vertical line at k = 3.95 Å-1 is the antinode
location of the lead only adsorbed on hematite and the vertical line at k = 3.45 Å-1
is the location of the Pb-Citrate antinode. The different lead concentrations were
used to maximize surface coverage of lead...... 148!
Figure 4.5. Results of the Linear Combination fit to the Pb EXAFS...... 149!
Figure 4.6. EXAFS data, fit and contribution from each path from (A) lead only and (B)
Lead and Fulvic acid on the HSA hematite demonstrating the impact the second
oxygen shell and two iron shells have on the overall shape of the EXAFS
spectrum...... 150! xviii Figure 4.7. XR (thick black lines) and RAXR results of lead and citric acid (red area) as
well as the lead only RAXR from chapter 3 (gray line) on the three hematite
surfaces...... 151!
Figure 4.8. EXAFS of phthalate and lead adsorbed on hematite at listed pH. Data in
blue, fit in red. The vertical line at k = 3.95 Å-1 is the antinode location of the
lead only adsorbed on hematite and the vertical line at k = 3.45 Å-1 is the location
of the Pb-Citrate antinode...... 152!
Figure 4.9. XR (thick black lines) and RAXR results of lead and phthalic acid (red area)
as well as the lead only RAXR from chapter 3 (gray line) on the three hematite
surfaces...... 153!
Figure 4.10. EXAFS of Fulvic acid and lead adsorbed on hematite at listed pH. Data in
blue, fit in red. The vertical line at k = 3.95 Å-1 is the antinode location of the
lead only adsorbed on hematite and the vertical line at k = 3.45 Å-1 is the location
of the Pb-Citrate antinode...... 154!
Figure 4.11. XR (thick black lines) and RAXR results of lead and fulvic acid (red area)
as well as the lead only RAXR from chapter 3 (gray line) on the three hematite
surfaces...... 155!
Figure 4.12. EXAFS of humic acid and lead adsorbed on hematite at listed pH. Data in
blue, fit in red. The vertical line at k = 3.95 Å-1 is the antinode location of the
lead only adsorbed on hematite and the vertical line at k = 3.45 Å-1 is the location
of the Pb-Citrate antinode...... 156!
xix Figure 4.13. XR (thick black lines) and RAXR results of lead and humic acid (red area)
as well as the lead only RAXR from chapter 3 (gray line) on the three hematite
surfaces...... 157!
Figure A.1. TEM images of (A) high surface area (99 m2/g) and (B) low surface area (35
m2/g) hematite ...... 170!
Figure A.2. Optimized structures for A) protonated and B) deprotonated outer-sphere
complex (OS-BN), C) protonated and D) deprotonated mononuclear bidentate
complex bound by the central and one terminal carboxyl (CT-MN), E) protonated
and F) deprotonated binuclear bidentate complex with the central and terminal
carboxyl groups (CT-BN), and G) protonated and H) deprotonated binuclear
bidentate complex bound by the central carboxyl and deprotonated hydroxyl
group (CH-BN). Atoms outlined in red are part of the iron oxide cluster, atoms
outlined in blue are the explicit water molecules and the atoms outlined in green
are the citrate molecule...... 171!
Figure A.3. Comparison of theoretical citrate spectra with and without (IEFPCM only)
explicit water molecules. Without explicit water molecules, the C-O-H bending
peak at1468 cm-1 is too intense and at too high of a wavenumber. Adding the
explicit water molecules results in a more realistic spectrum...... 172!
Figure A.4. Comparison of the theoretical asymmetric and symmetric stretch peak
locations with the experimental locations on the low surface area hematite at pH
2.5 and 7.5 ...... 173!
Figure A.5. Simplified representation of (A) Outer-sphere and (B) inner-sphere
complexes of citrate on hematite. The Inner-sphere complex does no occur on the xx atomically flat (001) hematite surface. The additional iron layer on the inner-
sphere complex represents the more corrugated structure of other common
hematite faces...... 174!
Figure B.1. Cross-section with the active surface on the top (Left) and overhead view
(Right) of (A) (001) surface, (B) (012) surface and (C) (110) surface views of the
ideal oxygen truncated hematite surfaces. The directions of the crystallographic
axes are also shown for reference. The surfaces vary in the proportions of singly,
doubly and triply coordinated oxygen atoms as well as the surface topography.186!
Figure B.2. Lead (0.1 mM) speciation in the a (0.1M NaClO4) background electrolyte
calculated by Visual Minteq3.0 ...... 187!
Figure B.3. Lead (0.1 mM) Speciation in the presence of 1 mM Citrate in a (0.1M
NaClO4) background electrolyte calculated by Visual Minteq 3.0...... 188!
Figure B.4. Lead (0.1 mM) Speciation in the presence of 1 mM Phthalate in a (0.1M
NaClO4) background electrolyte calculated by Visual Minteq 3.0 ...... 189!
Figure C.1. A) schematic of the sample cell used for XR and RAXR experiments from
(Bellucci et al., 2015) B) image of the (001) hematite sample mounted in the cell
on the diffractometer. The (012) and (110) crystal were 1 cm square rather than 1
cm by 3 cm...... 202!
Figure C.2. Image taken by the CCD detector of the reflection off the (110) hematite
surface. The two reflections are present due to the miscut in the crystal. The
overlap of the two rectangles in the area is integrated to get the reflection and the
area inside the rectangles, but outside of the overlap is used to determine the
background. The dark blue around the edges is the result of the slits placed before xxi the detector. These slits closed more when taking images around the Bragg peak.
If the reflection at the Bragg peak hits the detector, it could possibly damage the
detector...... 203!
Figure C.3. Example of the XR Data collected on the 110 surface. The top plot shows
the raw data with empty red circle. The green line is the ideal termination if no
adsorbents were present. The black line is the fit to the data. The solid blue
circles are the locations where the structure factor was taken for the RAXR fit.
The middle plot shows the data normalized to the generic CTR. The bottom plot
shows the residuals...... 204!
Figure C.4. Example of a model independent fit of the RAXR data on the (110) hematite
surface. The red circles show the data and the blue line is the fit. A baseline is
applied to both the data and the fit. In this method, each of the individual spectra
are fit and the lead location is determined by combining the fit...... 205!
Figure C.5. Example of a model dependent fit of the RAXR data on the (110) surface.
Red circles are the data and blue line is the fit. This is the same data set as in
Figure C3. For the model dependent fit, the fit lines are determined from the
model of lead locations. The best fit is then determined from the comparison of
the data to those fit lines...... 206!
Figure C.6. Comparison of the amplitude (top) and phase (bottom) of each RAXR
spectrum (e.g., data in Figures C.3 and C.4). The circles are from the model
independent fit (figure C.3) and the line is derived from the model dependent fit
(figure C.4)...... 207!
xxii Figure D.1. Model fit (blue line) to the RAXR data (red circles) for lead only at pH 6 on
the (001) surface...... 210!
Figure D.2. Model fit (blue line) to the RAXR data (red circles) for lead only at pH 6 on
the (012) surface...... 211!
Figure D.3. Model fit (blue line) to the RAXR data (red circles) for lead only at pH 6 on
the (110) surface...... 212!
xxiii
Chapter 1: Introduction
1.1 Problem Description
A recent National Research Council study puts the number of documented legacy contaminated waste sites across the country at over 126,000 with at least 12,000 of them being considered “complex” such that remediation is not likely to reduce contaminant levels below MCLs in the 50 – 100 year time frame (NRC, 2012). For years the ATSDR has placed lead second in overall concern on the CERCLA list of hazardous substances, behind only arsenic, and first in terms of occurrence at superfund sites (ATSDR, 2013).
Many of the superfund sites are due to centralized industrial activity, such as lead mining and processing, however decades of using lead as a gasoline additive and major component of paint have also resulted in an increased level of lead in nearly all soils
(Boutron et al., 1994). As paint and gasoline are used in higher concentrations in population centers, lead is found disproportionately in the soils of inner cities and other poverty stricken areas where people are least likely to be aware of the potential problem and least able to take precautionary measures to prevent harm from coming of it (Mielke et al., 1999; Mielke et al., 2011). The lead in these urban centers is also more bioavailable than lead pollution resulting from mining or smelting operations (Ryan et al.,
2004).
1 Virtually all bodily systems are negatively affected by elevated concentrations of lead
(Goyer, 1993), however it is especially potent as a neurotoxin which prevents the development of the brain and central nervous system in young children resulting in a permanent loss as measured by IQ as the child grows older (Lanphear et al., 2005). Due to lower IQ and an increase in antisocial behavior, lead has been linked to higher crime rates (Nevin, 2000). One of the major reasons for lead’s toxicity is that the body attempts to use lead in place of calcium, however while lead fits in the place of calcium, it does not provide the same functionality(Needleman, 2004). As a mimic of calcium, lead remains stored in bones for an extended period of time and can be released later in life when bone mass is reduced (Needleman, 2004) or during pregnancy or lactation, exposing the child to lead (Silbergeld et al., 1988). In addition to the high cost of lead pollution due to loss of life and health, there is also a high economic cost. The EPA has estimated the cost of cleanup for superfund sites to be in the range of $6-8 Billion per year for the next 25 years or up to $200 Billion (USEPA, 2004).
To clean up these sites efficiently and prevent lead from reaching drinking water sources requires an accurate estimate of lead speciation and fate. Transport and geochemical speciation models are created and applied to the contaminated sites to accomplish this goal. Unfortunately, lead transport does not always follow the behavior predicted of it in these models (Kaste et al., 2006). A primary reason for lead not following standard predictive transport models may lie in the oversimplification of models as they do not take into account the many heterogeneities in soil, including the role of mobile colloidal particles, size dependent oxide reactivity and natural organic matter (Hassellov and von
2 der Kammer, 2008; Tang and Weisbrod, 2009). Thus there is a critical need for research to elucidate the interactions of lead with mineral phases and organic acids.
1.1.1 Organic acid adsorption
Organic acids are ubiquitous in natural systems in the form of both low molecular weight acids with known structures, like citric acid and phthalic acid, and larger compounds with no set structure, such as humic and fulvic acids (Tan, 2011). The larger organic acids are present as a result of the breakdown of organic matter and are operationally defined.
Fulvic acid is soluble at all pH conditions whereas humic acid is insoluble below a pH of
2 (Sutton and Sposito, 2005). This is a result of the structure of the acids with humic acid having more phenolic functional groups and fulvic acid containing more carboxylic acid functional groups (Ritchie and Perdue, 2003). There is no consistent pKa for humic or fulvic acids as there is no set structure, but rather they become more negatively charged as the pH rises and the functional groups deprotonate (Stumm and Morgan, 1996).
Organic acids play many important roles in environmental systems such as acting as metal chelators in plants and microorganisms to aid nutrient uptake (Hell and Stephan,
2003) or sequestering harmful metals (Barone et al., 2008), acting as an easily accessible
“storage” for essential plant nutrients such as phosphate, and impacting the stability and reactivity of mineral particles. The focus of this dissertation is on how the organic acids influence the reactivity of the hematite surface with respect to the adsorption of lead.
The surface charge of hematite is pH dependent, being positive under acidic condition and negative under basic conditions with a pHpzc around 9, depending on the synthesis method (Rustad et al., 1999). When organic acids are included in the system, they have
3 the potential to adsorb on to the mineral surface, thus altering the surface charge.
Changing the surface charge will alter the adsorbent and the aggregation properties of the mineral particle (Stumm and Morgan, 1996). Organic acids are typically more attracted to iron oxides under mildly acid pH conditions where the positive surface will electrostatically attract the negatively charged acid. Under very acidic pH, below the acid’s pKas, the acid is neutrally charged and is not as likely to adsorb. This results in an adsorption peak in monoprotic acids and a plateau in multiprotic acids in the mildly acidic range (Stumm and Morgan, 1996).
The second chapter of this dissertation focuses on the adsorption of the low molecular weight organic acid, citrate on hematite and Table 1.1 lists a brief summary of recent papers focusing on citric acid adsorption onto various mineral phases. Citrate is an important compound used externally by plants and microorganisms to chelate iron and absorb it more efficiently (Yue et al., 2003) and sequester harmful metals. Internally, citric acid is part of the eponymous citric acid cycle which is critical for energy generation in all aerobic organisms (Madigan et al., 2008). Hematite and goethite are the most stable iron oxide and iron hydroxide respectively and are commonly found in natural soils (Cornell and Schwertmann, 2003). The two studies listed in Table 1.1 of citrate on the hematite surface are part of the same study. The results of these studies indicates that citrate is bound as an inner sphere complex, both singly protonated and fully deprotonated(Kallay and Matijevic, 1985; Zhang et al., 1985). More recently, several groups have investigated citric acid adsorption on goethite. No consensus on
4 exact adsorption modes has been released, but the overall view is that citrate adsorbs as both inner and outer-sphere complexes (Lindegren et al., 2009; Yeasmin et al., 2014).
1.1.2 Lead adsorption
Lead adsorption is an important process in understanding the mobility of lead in the environment. Mineral surfaces can act as either sinks of lead, immobilizing it (Bolan et al., 2014), or transport vectors, increasing the distance lead can travel (McCarthy and
Zachara, 1989). There are more reactive adsorbent minerals in natural systems, such as manganese oxides (O'Reilly and Hochella Jr, 2003), and preferable transformation processes, such as mineralization with phosphate (Ryan et al., 2004), however the large amount of iron oxides in soils makes them one of the most important adsorbent phases
(Cerqueira et al., 2011; Sauvé et al., 2000). Like most metals, lead has a steep adsorption edge on metal oxides. On hematite, the edge is located between pH of approximately 4 and 6 (McKenzie, 1980). At the surface of hematite, lead adsorbs as an inner-sphere bidentate complex (Bargar et al., 1997; Lenhart et al., 2001). Lead adsorbs strongly to hematite and it has been investigated as a potential sorbent for lead removal from drinking water (Shipley et al., 2013).
It is apparent from experiments where the direction of crystal growth is affected by the addition of organic acids that adsorption is not uniform on particles (Cho et al., 2009;
Cornell and Schwertmann, 2003). Mineral particles can contain several defined faces with different reactivities (Giammar et al., 2007; Venema et al., 1998). Hematite morphology can change dramatically depending on the synthesis methods leaving different crystal faces exposed (Schwertmann and Cornell, 2008). To shed more light on
5 the exact adsorption mechanisms, studies of lead on single crystals have been performed.
Table 1.2 lists some of the recent single crystal work performed using hematite, showing the variety of methods available.
Catalano et al. have performed X-ray reflectivity experiments on the three most common hematite faces, the (001) (Catalano, 2011), (012) (Catalano et al., 2007), and (110)
(Catalano et al., 2009). These three studies reveal that water forms a strongly ordered structure near the surface of the (012) and to a lesser extent the (110) surface while forming a very weakly ordered layer on the (001). This likely is related to the reactivity of the surface. The (001) surface is flat and uncharged in the environmentally relevant pH range (Venema et al., 1998) while the (012) and (110) have a more corrugated surface and carry a pH-depende surface charge (Shimizu and Boily, 2015). Bargar et al. (Bargar et al., 2004) performed grazing incidence extended X-ray adsorption fine edge spectroscopy of lead adsorbed on the (001) and (012) surfaces of hematite, finding both to be excellent sorbents for lead, though it was mostly adsorbed in the form of an oligomeric lead complex. Catalano et al (Catalano et al., 2006)also investigated the (100) surface, which is not one of the surfaces examined in this dissertation. In Catlalano’s study, selenite was reacted with the surface and using X-ray standing wave (XSW), they determined that the selenite adsorbed in a bidentate bridging manner to the surface through only the singly coordinated surface oxygen atoms. The doubly-coordinated surface oxygen atoms were less reactive. Furthermore, Se-Fe distances observed in an
XSW study may be misinterpreted in an EXAFS study as an edge sharing complex,
6 highlighting the importance of single crystal and EXAFS experiments being used together.
More recently, computational methods have been used to investigate both the hematite surface and the adsorption of lead to the surface (Table 1.2). Matching computational results with particle based experiments is often accomplished by using a small cluster of atoms to stand in for the particle. Although this is not a realistic scenario, using clusters has been shown to produce good results (Paul et al., 2007) and was used in chapter 2 of this dissertation to assist in interpreting the experimental infrared spectra. As single crystal methods use a specific surface, computational molecular modeling is a natural complimentary method as the exact structure of the sorbent surface is input by the experimenter. Trainor et al. (Trainor et al., 2004) combined nonresonant X-ray reflectivity (XR) and density functional theory (DFT) modeling to study the (001) surface of hematite finding that the surface that was most likely exposed was not ideally terminated but was partially a half layer termination resulting in singly coordinated oxygen in addition to the doubly coordinate oxygen of the ideal termination. Kerisit
(Kerisit, 2011) modeled all three hematite surfaces that Catalano investigated with molecular dynamics, finding good agreement between the theoretical and experimental water binding. Triply coordinated surface oxygen form much stronger hydrogen bonds than do the doubly coordinated oxygen atoms dominating the (001) surface explaining the strong water ordering of the (012) and (110) and weak water ordering of the (001)
(Catalano et al., 2007; Catalano, 2011; Catalano et al., 2009; Kerisit, 2011).
7 Recently, the theoretical work has begun to outpace the experimental work. Mason et al.
(Mason et al., 2009) investigated the adsorption of Pb to the (001) surface of hematite.
Mason found that the lead adsorbs on the hematite surface as either a bi- or tridentate complexes on the surface with Pb-Fe distances ranging from 2.21 to 2.34 Å. The particle based experiments of lead yield a distance of ~2.3 Å, within the computational margins.
GIEXAFS gives the Pb-Fe distance on the (001) surface as 2.24 Å (Bargar et al., 2004).
1.1.3 Interaction of metals and organic acids
Combining metal cations and organic acid and other competing anions together in systems with sorbent surfaces moves the test systems closer to what is observed in the environment. There has been considerable research on the topic. Table 1.3 summarizes the recent work done on lead and anions together on mineral surfaces. Generally speaking, the presence of organic acids tend to increase the amount of lead adsorbed on oxide surfaces under low pH conditions and hinders the adsorption of lead at a higher pH.
There are three types of ternary complexes reported in the literature for metal – anion binding to mineral surfaces, metal bridging, outer-sphere metal acid complex and acid bridging / film. Lead has been found adsorbed directly to the mineral surface forming a bridge to the organic acid which is bound to the adsorbed lead. Malonate (Lenhart et al.,
2001) has been observed in this arrangement on hematite, as has sulfate (Ostergren et al.,
2000). Humic acid (Orsetti et al., 2006) and sulfate (Swedlund et al., 2009) were observed to form the same ternary structure on goethite. Bargar et al. found EDTA complexed lead and then formed an outer-sphere PbEDTA2- complex on goethite (Bargar et al., 1999). Yip et al, found the same Pb-EDTA outer-sphere complex formed on
8 goethite when the goethite was coated with lead before the EDTA was added to the system, whereupon the EDTA removed the lead from the surface and adsorbed as an outer-sphere complex (Yip et al., 2010). Simanova (Simanova et al., 2011) found that cobalt and oxalate form an outer-sphere complex that transitions over time into a bridging inner sphere complex. The third adsorption structure involves the metal to be adsorbing in a film. Templeton et al. (Templeton et al., 2001) investigated lead adsorption on the hematite (001) surface pre-equilibrated with a biofilm. They found that under low lead loading conditions that lead adsorbed directly to the hematite surface, however, as the lead concentration was increased, the lead interacted more with the biofilm and becoming an important lead sink. This adsorption structure has also been reported by Lee et al. where lead adsorbed on the muscovite surface in a fulvic acid film (Lee et al., 2011).
1.2 Research objectives
This work is focused on elucidating the interactions of organic acids and / or lead at the hematite water interface. A greater understanding of the behavior of lead and organic acids will increase our ability to predict the fate of lead in the environment. There are four research objectives in this dissertation:
1) Determine the bonding mode of citric acid on hematite nanoparticles.
2) Further the knowledge of lead adsorption on hematite with single crystal studies.
3) Determine the effect of organic acids on the adsorption of lead to hematite on
both nanoparticles and specific crystal faces.
4) Merge the results of the single crystal work with the particle based studies.
9 1.3 Dissertation overview
This dissertation has three main Chapters. First studying the adsorption of an acid, citric acid, on hematite particles (Chapter 2). Second using several synchrotron based techniques to further investigate the adsorption of lead on hematite (Chapter 3). In the third, we studied lead adsorption on hematite in the presence of four organic acids
(Chapter 4). Chapter 5 contains a summary of the work and suggestions for future research.
1.3.1 The impact of particle size on the adsorption of citrate to hematite
The second chapter of this dissertation deals with how citrate adsorbs on the surface of hematite particles. This was accomplished using batch adsorption, infrared spectroscopy, density functional theory molecular modeling and surface complexation modeling.
Citrate on hematite has not been studied as much as citrate adsorption on the goethite in recent years as evidenced by the summary in Table 1.1. The study was performed with two different sized hematite nanoparticles to determine if there is any size dependent reactivity. In this chapter we hypothesized that we the citrate would be reactive with the hematite surface and the reactivity would vary with size. This chapter is in press at the
Journal of Colloid and Interface Science with co-author John Lenhart.
1.3.2 X-ray analysis of lead adsorbed on the hematite (001), (012), and (110) surface
The third chapter uses two synchrotron based methods, extended X-ray absorption fine structure (EXAFS), and X-ray reflectivity, to investigate the mechanisms of lead adsorption on hematite. EXAFS is an element specific method to determine the local environment around the target atom and it was performed on lead adsorbed to hematite 10 particles. The X-ray reflectivity was performed on three different single crystal surfaces, the (001), (012), and (110). We hypothesized that the lead would adsorb as an inner- sphere complex, however the exact mechanism and amounts of adsorption would vary with the hematite face. This chapter was written to be submitted to Environmental
Science and Technology and was co-authored by John Lenhart and Sang Soo Lee at
Argonne National Lab.
1.3.3 Lead and organic acids on hematite
The impacts on the surface structure of lead in the presence of organic acids was investigated by combing lead with citric, phthalic, humic and fulvic acids in turn. The same techniques and surfaces used in chapter 3 where employed here. We hypothesized that the organic acids would impact the quantity of lead adsorbed and the manner in which it is adsorbed to the surface. Again we thought the individual surfaces would play an important role in the adsorption modes. This chapter was written for submittal to
Geochimica et Cosmochimica Acta with co-authors John Lenhart and Sang Soo Lee.
11
Table 1.1. Summary of recent / relevant studies on the adsorption of citrate on mineral surfaces.
Sorbent Methoda Result Source Citrate adsorbed as inner sphere protonated and (Zhang et al., Hematite Dissolution deportonated complexes 1985) Zeta potential & Citrate bound as a bidentate structures as both singly (Kallay and Hematite SCM and fully deprotonated. Matijevic, 1985) Citrate adsorbs via an inner-sphere bidentate (Hidber et al., Corundum ATR-FTIR complex 1996) Inner-sphere at low pH and outer-sphere at high pH. Goethite, illite, ATR-FTIR / Citrate adsorbed much more to goethite than illite or (Lackovic et al., kaolinite BA kaolinite 2003) Citric acid adsorbs directly to the surface through binding with the calcium and hydrogen bonding with the oxygen groups. The citrate was bound as a Molecular bridging bidentate complex where surface geometry (Filgueiras et al., Apatite Simulations allowed. 2006) Citrate was photooxidized to acetonedicarboxylic acid when sorbed to the surface. In the presence of excess citrate, the newly formed acetonedicarboxylic acid was replaced at the surface by citrate through (Borer et al., Lepidocrocite ATR-FTIR ligand exchange reaction. 2007) Molecular Citric acid adsorbs preferentially to the (0110) (de Leeuw and Hydroxyapatite dynamics surface over the (0001) surface. Rabone, 2007) Multiple inner and outer-sphere complexes exist. Evidence was found for the hydroxyl group (Lindegren et al., Goethite ATR-FTIR deprotonating and playing a role in the adsorption 2009) Citric acid adsorbs preferentially to the (0001) ZnO face in an inner sphere manner causing growth on the ZnO Crystal growth (0001) face to be suppressed. (Cho et al., 2009) Inner sphere deprotonated citrate is present on the surface of TiO2 across the range of pH tested (Mudunkotuwa indicating the pKas for citrate are reduced at the and Grassian, TiO2 ATR-FTIR surface. 2010) Citrate adsorbs as both mono- and bi-dentate inner sphere complexes as well as an outer-sphere complex. The citrate was found to have a much Goethite, clay higher affinity to the goethite than the clays tested, (Yeasmin et al., mineral ATR-FTIR /BA kaolinite, illite or montmorillonite 2014) a SCM = Surface Complexation modeling; ATR-FTIR = attenuate total reflectance Fourier transform infrared spectroscopy; BA = batch adsorption
12
Table 1.2. Summary of single crystal hematite experiments.
Surface Adsorbate Methoda Results Source Lead adsorbs directly to the hematite surface in the presence of a biofilm at low Pb concentrations. As Pb + the Pb concentration increases, the biofilm becomes a (Templeton et (001) Biofilm XSW more important sink. al., 2001) (001), Oligomeric lead compound formed at both surfaces (Bargar et al., (012) Pb2+ GIEXAFS with surface coverage above 2 µmol/m2 2004) Reported the presence of 2 distinct terminations (Trainor et al., (001) XR under a near water saturated helium atmosphere. 2004) U(VI) adsorbs on the corundum surface as a monodentate complex, but on the hematite surface as a bidentate complex indicating that the surface XR / structure alone does not determine the adsorption (Catalano et (012) U (VI) GIEXAFS mode. al., 2005) SeO3 adsorbs as a bridging bidentate complex between adjacent singly coordinated oxygen atoms, not with the doubly coordinated oxygen, and the Se- Fe distance could be mistaken for an edge-sharing (Catalano et (100) SeO3 XSW complex in EXAFS al., 2006) Strong interfacial ordering following the topography (Catalano et (012) Water XR of the crystal al., 2007) The termination of the (012) surface changes with (Lo et al., (012) Water DFT temperature, therefore thermal annealing is critical 2007) The (012) surface shifts between full and half layer terminations depending on preparation. Thermal (Tanwar et al., (012) -- XR annealing is critical to achieve a full termination. 2007) There is simultaneous inner and outer sphere (Catalano et (012) AsO4 XR / RAXR adsorption al., 2008) Strong interfacial ordering following the topography (Catalano et (110) Water XR of the crystal al., 2009) Lead adsorbs more strongly to hematite than (Mason et al., (001) Lead DFT Corundum. 2009) There is weak layering of interfacial water on the (Catalano, (001) Water XR (001) surface 2011) (001), There is greater charge storage capability on the (Shimizu and (012) -- EIS (012) face than the (001) face Boily, 2015) Two terminations were found on the hematite (001) XR / ζ surface, but the potential indicates that the surface is (Lutzenkirchen (001) water potential still dominated by doubly coordinated oxygen. et al., 2015) Fe(II) can adsorb on the (001) surface as either a mono-or tri-dentate inner-sphere complex or an outer (Kerisit et al., (001) Fe(II) MD sphere complex. 2015) a (XSW = X-ray standing wave, GIEXAFS = grazing incidence extended X-ray fine structure, XR = nonresonant X-ray reflectivity, RAXR = resonant anomalous X-ray Reflectivity, DFT = density functional theory, EIS = electrochemical impedance spectroscopy, MD = molecular dynamics)
13 Table 1.3. Summary of recent studies on the adsorption of lead to mineral surfaces in the presence of organic aids or other anions. Anion Mineral Results Source Lead and EDTA form a PbEDTA2- complex that is held to EDTA Goethite the goethite surface as an outer-sphere complex (Bargar et al., 1999) Pb adsorption is enhanced in the presence of SO4. The SO4 causes lead to adsorb in a corner-sharing manner due to the formation of a ternary complex with the SO4 having a Sulfate Goethite stabilizing effect on the sorbed lead (Ostergren et al., 2000) Pb adsorbs directly to the hematite surface and malonate Malonate Hematite forms a ternary complex by complexing bound lead (Lenhart et al., 2001) EDTA, Citric acid Goethite EDTA decreases the amount of lead adsorbed to goethite (Wu et al., 2003) Fulvic acid increased the amount of lead adsorbed across the (Heidmann et al., Fulvic Kaolinite range of pH, especially under lower pH conditions 2005) Under the acid pH conditions tested, Pb forms a bridge Humic Goethite between the humic acid and the surface (Orsetti et al., 2006) Surface complexation modeling was performed to verify the (Swedlund et al., Sulfate Goethite ternary Pb-SO4 complex 2009) EDTA removes inner-spherically bound lead from the goethite surface replaced by a Pb-EDTA outer-sphere EDTA Goethite complex. (Yip et al., 2010) Citric, The adsorption edge is shifted to the lower pH range in the (Perelomov et al., Oxalic Goethite presence of citric and oxalic acids. 2011)
14
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22
Chapter 2: The Impact of Particle Size on the Adsorption of Citrate to Hematite
(Journal of Colloid and Interface Science, In Press)
Abstract
We investigated the adsorption of citric acid on the surface of two different sized hematite nanoparticles using batch adsorption experiments, Fourier-transform infrared
(FTIR) spectroscopy, surface complexation modeling and computational molecular modeling. Citrate adsorption was at a maximum between pH approximately 2.5 and 5.5 and declined as the pH was increased or decreased from that range. FTIR analysis and computational molecular modeling indicated the dominant adsorption mode across the range of pH tested was an outer-sphere complex when the citrate concentration was high enough to saturate the surface. At low pH, there was also evidence of an inner-sphere binuclear, bidentate complex where the hydroxyl group was deprotonated and played an active role in the adsorption. The structure of the outer-sphere complex slightly varied with pH, existing in the singly protonated state at low pH and the fully deprotonated state at high pH. Although the outer-sphere complex was dominant in terms of total surface coverage when the surface was saturated with citrate, the inner-sphere complex appears to be the dominant adsorption mode in lower citrate concentrations. Surface area- normalized surface coverages were similar for both sizes of hematite, however, slight differences were observed in the adsorption tendencies of citrate to the different sized
23 particles. The inner sphere complex was more prevalent on the smaller particles than the larger particles. The adsorption of citrate on hematite was described with a triple layer model using one inner-sphere complex with the hydroxyl and all three carboxyl groups deprotonated, and two outer-sphere species, one fully deprotonated and the other singly protonated. The determined equilibrium constants were applied to a number of experimental systems and returned adequate results. The results indicate a minor dependence in citrate adsorption on particle size, however a greater variety of particle sizes and morphologies is needed to draw more firm conclusions.
2.1 Introduction
Organic acid interactions with metal oxide surfaces are a ubiquitous presence in natural soils and play an important role in many biological and geochemical processes (Axe et al., 2006). When bound to an oxide surface, organic acids can alter the surface properties of the underlying mineral, which modifies its solubility and interactions with other solutes (Zinder et al., 1986). Organic acids also influence particle fate in the environment by altering the steric and electrostatic interactions that control aggregation and transport
(Davis, 1982; Hassellov and von der Kammer, 2008; Zinder et al., 1986). In addition to the effect organic acids have on particle surface reactivity they also bind metals in solution, either preventing or enhancing the metal’s adsorption (Davis and Leckie, 1978) and thus organic acids can impact metal transport through soil (Hassellov and von der
Kammer, 2008; Kaste et al., 2006; McCarthy and Zachara, 1989; Teutsch et al., 2001).
Therefore, understanding the conditions that control organic acid adsorption to metal
24 oxide surfaces and the coordination modes of those acids will lead to a better understanding of metal transport in contaminated soils.
There are a wide variety of organic acids present in natural soils that can bind to metal oxides. Humic and fulvic acids are among the most prevalent and they are comprised of large polyelectrolytic, polyfunctional acids heterogeneous in size and structure (Cabaniss et al., 2000; Lenhart et al., 2000). While these humic substances are ubiquitous in soils and play an important role in soil chemistry, elucidating their fundamental adsorption mechanisms remains a challenge (Tan, 2011). One alternative is to instead use low molecular weight organic acids of known structure that share common functional groups
(e.g., carboxylic acid) with heterogeneous humic substances (Evanko and Dzombak,
1998; Lenhart et al., 2000). In this study, citric acid was used as a representative organic acid due to its prevalence in the environment and functional group similarities with humic substances (Strobel, 2001). Citric acid is used by plants and microorganisms as a metal chelator in iron and phosphate deprived soils to more efficiently extract the iron and phosphate necessary for growth (Hell and Stephan, 2003; Pierre and Gautier-Luneau,
2000) and in low pH soils to prevent aluminum toxicity (Barone et al., 2008). Citric acid has three carboxylic acid functional groups, with intrinsic pKas of 3.13, 4.76, and 6.40,
(Stumm and Morgan, 1996) and one hydroxyl group that deprotonates at pH values estimated to be between 11 and 14.4 (Silva et al., 2009). The pK value for the hydroxyl group is out of the range of this study; however, that does not preclude it from deprotonating and playing an active role in the adsorption process as it does when complexing aluminum or gallium (Clausén et al., 2005).
25 Hematite (α-Fe2O3) is thermodynamically very stable and it is a common end result of the transformation of other less stable iron oxides (Cornell and Schwertmann, 2003b). As such it is commonly found in natural soils, especially older soils, in both nanoscale and larger forms (Cornell and Schwertmann, 2003b). Hematite readily binds organic and inorganic species and in a colloidal form it is implicated in the larger than anticipated transport of heavy metals through sediment and subsurface systems (Catalano et al.,
2008; Gimenez et al., 2007). The different crystalline faces of hematite display varying reactivity and available binding sites. For example, the (001) face, considered the most common face, has a net neutral surface charge between pH 2 and 10, (Hiemstra and Van
Riemsdijk, 1999) while the pHpzc of whole hematite particles is reported to be between
8.5 and 9.5 (Hwang and Lenhart, 2008; Rustad et al., 1999). This lack of surface charge in the environmentally relevant pH range results in the (001) surface being less reactive than other common surfaces and makes aggregation and adsorption processes on the
(001) surface less dependent on pH (Hiemstra and Van Riemsdijk, 1999).
The size of the hematite particles is expected to play an important role in adsorption processes as the relative abundance of the different crystalline faces and their associated binding sites may change with particle size (Gaboriaud and Ehrhardt, 2003; Madden et al., 2006). For example, Madden et al. (Madden et al., 2006) determined that the affinity of Cu(II) to hematite changed with particle size as the smaller hematite had a higher proportion of irregular octahedral binding sites preferred by Cu(II). This occurred even though both particles exhibited the same pseudo hexagonal morphology. Similar results were also demonstrated with the iron hydroxide goethite, which displays a better defined
26 crystal morphology than does hematite (Gaboriaud and Ehrhardt, 2003). To isolate the role of different crystalline faces of goethite, atomic force microscopy was used to determine the (001) face made up 70% of the surface of larger crystals versus 30% of that for smaller crystals (Gaboriaud and Ehrhardt, 2003). Thus, the adsorption capacity of the goethite crystals will not scale directly with available surface area if the adsorbate has a preference for sites in the (001) face (Gaboriaud and Ehrhardt, 2003). In this study we use two different sizes of hematite to investigate the effect of particle size, and by inference, crystal face distribution on the adsorption of citrate to hematite.
In comparison to the extensive literature of citrate adsorption on goethite and other oxides, details of citrate adsorption on hematite are limited. For example, Zhang et al.
(Zhang et al., 1985) used macroscopic adsorption and iron dissolution experiments to conclude that citrate adsorbed as both doubly and triply deprotonated species. As part of the same study, Kallay et al. (Kallay and Matijevic, 1985) used zeta potential measurements and surface complexation modeling to determine that citrate is bound directly to the hematite surface via a bidentate structure and that both singly protonated and fully deprotonated citrate surface species exist. Additional research has been performed on the iron oxyhydroxide, goethite, and the isostructural aluminum oxide, corundum. For example, Hidber et al. (Hidber et al., 1996) used a variety of methods, including Fourier transform infrared (FTIR) spectroscopy, to investigate citric acid adsorption on corundum (α-Al2O3). They reported citrate adsorbs in an inner sphere manner, but not with all three carboxyl groups. Hidber et al. (Hidber et al., 1996) did not find direct evidence for the involvement of the hydroxyl group in citrate binding but
27 assume it is involved due to differences between its adsorption and that for tricarballylate, which differs from citrate by lacking the additional hydroxyl group. Results with goethite are more extensive, with early investigators primarily invoking inner-sphere surface complex modes, such as one involving the citrate ion adsorbing in a triply coordinated manner using all three carboxyl groups (Cornell and Schindler, 1980). More recent results have increasingly found outer-sphere adsorption modes, with Lackovic et al. (Lackovic et al., 2003) indicating inner-sphere complexation at low- to mid-pH gives way to outer-sphere complexation at elevated pH. This interpretation was based on FTIR spectroscopy coupled to batch adsorption and goethite dissolution experiments (Lackovic et al., 2003). Lindegren et al. (Lindegren et al., 2009) reported finding multiple inner- and outer-sphere surface structures using 2-D correlation infrared spectroscopy, including two previously unreported structures. These new structures included a protonated outer- sphere complex at low pH and an inner-sphere structure at high pH where the hydroxyl group deprotonated and played an active role in the adsorption process. Most recently,
Yeasmin et al. (Yeasmin et al., 2014) used batch adsorption experiments with 14C-labeled organic acids and infrared spectroscopy and came to a similar conclusion as Lindegren et al. that citrate adsorbs on goethite (and ferrihydrate) as mono- and bidentate inner sphere complexes as well as outer-sphere complexes. Unfortunately, they did not further specify the structure of the adsorbed citrate.
The main method of investigation in this study to determine the structure of adsorbed citrate onto hematite was Attenuated Total Reflectance FTIR (ATR-FTIR) spectroscopy.
As the different structural features of the citrate molecule absorb infrared energy at
28 specific frequencies, the changes in frequency and magnitude of that absorption can indicate changes in the coordination of the molecule (Hwang et al., 2007; Lackovic et al.,
2003; Lindegren et al., 2009). These experimentally determined infrared spectra were compared with those produced through computational methods. Calculating the theoretical infrared spectra of different potential structures allows us to assign specific vibrations to the peaks visible in the experimental FTIR spectra and compare the experimental results of the adsorbed structure to those generated theoretically based upon different potential structures (Hwang et al., 2007). Based on these constraints, a surface complexation model was derived based on the Triple Layer Model of hematite developed by Hwang and Lenhart (Hwang and Lenhart, 2008) and utilized to simulate the experimental adsorption data. Our results suggest that citrate adsorbs predominantly as an outer-sphere complex on both sizes of hematite studied, with a minor addition of a bidentate inner-sphere complex that adsorbs preferentially with respect to the outer- sphere complex, but in lower concentrations due to a limited number of available surface sites. The smaller hematite particles allowed for more inner sphere adsorption suggesting it carried a greater proportion of sites that prefer to directly bind citrate.
2.2 Materials and Methods
2.2.1 Hematite Synthesis and Characterization
Two sizes of hematite were synthesized and used in this research. The first, with a nominal 10 nm diameter was synthesized following methods in Madden et al. (Madden et al., 2006) by slowly dripping 50 mL of a 1 M solution of Fe(NO3)3 into 625 mL of
29 boiling Millipore water. The solution was allowed to cool gradually overnight before being cleaned. The second, with an average diameter of 50 nm, was synthesized through a forced hydrolysis method using ferric chloride as the iron source following the methods of Matijevik and Scheiner (Matijevic and Scheiner, 1978) and Penners and Koopal,
(Penners and Koopal, 1986) with minor modifications (Hwang and Lenhart, 2009).
Sufficient ferric chloride was added to 0.004 M HCl to create a 50 mL 0.8 M FeCl3 solution. This solution was filtered through a 0.22 µm PVDF filter and added to 1950 mL of 0.004 M HCl preheated to 98oC. This suspension was aged at 98 oC for 3 days before being rapidly cooled and cleaned. Both hematite suspensions were concentrated by adding solid NaOH to raise the pH leading to particle aggregation and settling. The concentrated hematite solids were collected and transferred to 8-10 kDa cellulose ester dialysis tubing (Spectrum Labs) where they were dialysed against Millipore water for several days. The dialysis water was changed twice daily until the dialysate conductivity approached that for deionized water. The suspensions were subsequently dialysed against a 1mM HClO4 solution for 24 hours. Multiple batches were made and combined to create stock solutions. For hematite synthesis and all following experiments, deionized water used with a resistivity of 18.2 µS/cm was supplied using a Millipore Milli-Q Plus system. All reagents were ACS grade or higher and the glassware was cleaned using a
5% nitric acid bath followed by repeated rinsing in deionized water. The hematite particle samples were verified as crystalline hematite using X-ray diffraction (either a
Scintag Pad V or Rigaku SmartLab). Particle size and morphology were determined using a Tecnai BioTwin TEM (see Figure A.1 in Appendix A). To image the particles, an aliquot of the stock hematite solution was diluted in deionized water and a drop was 30 placed on wax paper. A copper carbon formvar TEM grid was placed on the drop and allowed to rest for ~20 seconds before wicking the solution off of the grid and drying the sample under a nitrogen stream. Surface area was measured with gas adsorption using a
Micromeritics Flowsorb II 2300 BET instrument. The 10 nm particles had a BET surface area of 99 m2/g and the 50 nm diameter hematite had a surface area of 35 m2/g. Hereafter these particles are referred to as high and low surface area hematite (HSA and LSA), respectively.
2.2.2 Batch Adsorption Protocol
Samples of the hematite stock solutions were diluted to 10 g/L in 1mM HCl and purged overnight using humidified nitrogen gas to remove carbon dioxide. Aliquots of the CO2- free hematite were transferred to 50 mL polycarbonate centrifuge tubes. The appropriate amount of citric acid, NaCl and either HCl or NaOH were added, under a stream of humidified nitrogen gas, to achieve the desired acid concentration, background electrolyte concentration (0.1 M unless otherwise mentioned) and pH with a 5 g/L hematite concentration. While the type of background electrolyte can impact adsorption processes, (Criscenti and Sverjensky, 1999) NaCl was chosen because it is not infrared active, unlike other common electrolytes, NO3 and ClO4. All solutions were made using
CO2-free deionized water. Although Zhang et al. (Zhang et al., 1985) reported citrate adsorption was complete within 5 minutes, the sample tubes were allowed to equilibrate on an end-over-end rotator for 48 hours to ensure both adsorption and coordination equilibrium were reached. This equilibration process was conducted in the dark to avoid light-catalyzed side reactions (Borer et al., 2007; Dodge and Francis, 2002). After 31 equilibration, the final pH of the samples was measured and the samples were centrifuged for 30 minutes at 12000 rpm. The supernatant was analyzed for free citrate using a
Dionex ICS-2100 ion chromatography system immediately after centrifugation.
Dissolved iron was measured using an inductively coupled plasma – atomic emission spectrometer (Varian Vista AX CCD-Simultaneous ICP-AES). Minimal iron was found at pH 1.5 and none was measured at higher pH values, which was in agreement with previously published studies (Persson and Axe, 2005; Zhang et al., 1985).
2.2.3 ATR-FTIR Spectroscopy
ATR-FTIR spectroscopy was performed with a Thermo Nicolet Nexus 670 spectrometer using a duraSampl IR 9 bounce diamond coated ATR cell. The empty cell was used for the background spectra and a new background was collected before each sample. All sample and background spectra were collected by averaging a minimum of 256 scans, at a resolution of 4 cm-1, in Nicolet’s OMNIC software (v. 8). Reference spectra of known structures were collected to compare to the sorbed citrate spectra. Citrate spectra in the
NaCl background electrolyte were collected across a range of pH values to determine the baseline position and pH-dependence of the important citrate and citric acid IR-active groups. Ferric citrate spectra, created using a mixture of 0.2 M sodium citrate and 0.07
M FeCl3, were also collected across a range of pH values to serve as a model for inner sphere complexation. A 0.1 M NaCl solution was subtracted from all of the reference spectra to remove the influence of the H-O-H bending motion from water.
32 Sorbed citrate spectra were measured on the wet paste collected after centrifugal separation of the hematite at the conclusion of the batch sorption experiments. The paste was applied to the surface of the ATR crystal and covered with reserved supernatant and a crystal cap to prevent dehydration of the paste during measurement. In order to remove the dominant H-O-H bending vibration of water (at ~1635 cm-1), a spectrum from a hematite-only wet paste was subtracted from the citrate coated hematite sample spectra following the approach described in Hwang et al. (Hwang et al., 2007) and similar to the methods of Kubicki et al. (Kubicki et al., 1999; Kubicki et al., 1997) Subtracting the hematite wet paste spectrum rather than plain water or a supernatant spectrum was necessary to account for changes in the H-O-H intensity due to the physisorbed water at the surface of the hematite (Axe et al., 2006).
2.2.4 Molecular Modeling
Using a computational method to calculate theoretical vibrational spectra allowed us to test various proposed surface structures implied from the experimental IR spectra. A variety of plausible structures, both inner- and outer-sphere, were tested. One of the challenges in working with hematite is the difficulty in identifying specific crystalline faces (Cornell and Schwertmann, 2003a), making it challenging to perform theoretical analyses using a slab model. Thus, the hematite surface was modeled using a small cluster comprised of one or two iron atoms as this approach has been shown to perform well vs. periodic slab models (Paul et al., 2007). The makeup of the two types of clusters used were Fe(OH2)6 for mononuclear complexes and Fe2(OH)4(OH2)6 for binuclear complexes (Hwang et al., 2007; Kubicki et al., 2007). In the case of the inner sphere 33 complexes, the carboxylic groups from the citrate displaced OH2 groups. Both the aqueous and adsorbed citrate structures were optimized using density function theory
(DFT) at the B3LYP (Becke, 1993; Lee et al., 1988; Stephens et al., 1994) level of theory with 6-31+G(d) basis set for all atoms (Hehre et al., 1986). This same level of theory was used for the vibrational calculations. Solvation effects were taken into account by using the Integrated Equation Formalism – Polarizable Continuum Model (IEF-PCM) for water (Tomasi et al., 2005) as well as explicit water molecules (Hwang et al., 2007).
Calculations were performed with and without explicit water molecules. The addition of explicit water molecules was found to give superior results for the citrate only spectra, so only these results are reported. All calculations were performed using Gaussian 09
(Frisch et al., 2009) at the Ohio Supercomputer Center. Optimized structures are presented in Figure A.2. in Appendix A.
2.2.5 Surface Complexation Modeling
Surface complexation modeling was applied to simulate the results of the batch adsorption experiments to further validate the structures taken from the infrared and molecular modeling calculations by extending them to fit macroscopic trends in adsorption across a range in solution conditions. We chose to utilize the triple layer model (TLM) because it possesses the complexity necessary to simulate dual-mode adsorption, which previous research for citrate adsorption on goethite suggests is possible, (Lackovic et al., 2003; Lindegren et al., 2009; Yeasmin et al., 2014) and because TLM parameters for hematite used in this research were determined by Hwang and Lenhart (Hwang and Lenhart, 2008). These model parameters are summarized in 34 Table 2.1. All fits were done in FITEQL v. 4 (Westall and Chemistry, 1982). The goodness of fit was determined using a weighted sum of squares / degrees of freedom
(WSOS/DF) as determined by FITEQL. All of the fits utilized a relative error of 0.1 for pH and an absolute error for citrate of 0.01 * the total citrate concentration (Dzombak and
Morel, 1990; Katz and Hayes, 1995).
2.3 Results and Discussion
2.3.1 Adsorption data
The adsorption envelope for citrate on hematite (Figure 2.1) agrees with the expected adsorption of a multiprotic organic acid on an iron oxide surface with a broad adsorption maximum in the acidic to near neutral pH range and declining adsorption with increasing pH (Stumm and Morgan, 1996). As the pH increases, the charge on the citrate molecule changes from neutral below the first pKa of 3.13 to negative three above the final pKa of
6.40. The surface charge of the hematite also changes with pH, going from positive in the acidic pH region to zero at the pHpzc around pH 9 (Hwang and Lenhart, 2008) and then negative at higher pH values. The adsorption data shown for the different sized particles reflect the concentrations of citric acid necessary to saturate the particle surface.
While the high surface area hematite adsorbed over twice as much citrate when normalized to mass (Figure 2.1a), the maximum adsorption for the two particles was very similar when normalized to surface area (Figure 2.1b). The envelope shape, however, was subtly different for the two hematite types. For example, the HSA hematite featured a sharp adsorption maximum centered at pH 2.5 while the LSA hematite exhibited a
35 flatter peak between pH 2.5 and 5.5. At elevated pH values, the pH dependence of citrate adsorption to the HSA hematite was lower than that for the LSA hematite. The position of the adsorption edge for LSA and HSA coincided quite closely. This was in contrast to
Madden et al. (Madden et al., 2006) observing that the adsorption edge for Cu(II) shifted down 0.6 pH units on HSA hematite versus LSA hematite.
The change in the surface coverage of citrate on LSA hematite was evaluated at 0, 100, and 500 mM NaCl (Figure 2.2). At pH 2, an increase in NaCl from 0 to 500 mM reduced adsorption as the surface coverage of citrate at 500 mM NaCl was approximately 70 percent of that observed in the system with no extra NaCl. Ionic strength dependence like this is typically a marker of outer sphere adsorption as the background electrolyte can competitively adsorb in place of the citrate (McBride, 1997). As the pH increased, the charge on the hematite surface became less positive resulting in a lessening of the ionic strength dependence and at ca. pH 8 very little change in the adsorption of citrate was observed. Above the pHpzc approximately 20% of the maximum amount of citrate was still adsorbed (Figure 2.1). At this pH, both citrate and hematite were negatively charged making the formation of electrostatically bound complexes unlikely, although not out of the question. For example, Hwang et al. (Hwang et al., 2007) reported phthalic acid adsorbing predominantly as an outer sphere species on hematite at pH approaching and above the pHpzc. Persson et al. (Persson et al., 1998) found a similar pattern for phthalate on goethite where the phthalate adsorbed in an inner-sphere manner at low pH and outer- sphere at high pH. Likewise, Lackovic et al. (Lackovic et al., 2003) found that citrate
36 adsorbed on goethite predominantly as an inner-sphere complex, however, an outer- sphere complex was needed at high pH to adequately explain the adsorption behavior.
2.3.2 FTIR spectroscopy
2.3.2.1 Reference spectra
Infrared spectra of aqueous citrate and ferric citrate were collected across a range of pH values for use as reference spectra (Figure 2.3). The citrate spectra show a dominant
-1 peak at ~1720 cm at low pH, resulting from the carbonyl stretch (νC=O) associated with protonated carboxyl groups, which was reduced in intensity with increasing pH as the carboxyl groups deprotonated until it was absent at a pH of 5.5. As the intensity of the carbonyl peak waned with increasing pH, peaks at 1570 cm-1 and 1390 cm-1 appeared and correspondingly grew in intensity. These two features represent the asymmetric (νasC-O) and symmetric (νsC-O) carbon – oxygen stretches of carboxylate groups, respectively
(Lackovic et al., 2003; Lindegren et al., 2009). The peak at ~1226 cm-1 was also dominant at low pH and with increasing pH it decreased in conjunction with the carbonyl stretch suggesting it originated from protonated carboxylic acid groups. DFT calculations show this feature was related to C-O-H bending (δC-OH) of the carboxylic acid groups. The feature centered at about 1280 cm-1 can be resolved into three individual peaks at 1257, 1280 and 1296 cm-1 and from DFT analyses the peak at ~1280 cm-1 was related to the C-O-H bending of the hydroxyl group while the two peaks surrounding it resulted from C-H rocking motions from the citrate skeleton. From these simulations, the shoulder in the symmetric stretch at ~1435 cm-1 resulted from both the
37 C-H rocking motions on the citrate skeleton as well as the C-O-H bending motion of the hydroxyl group.
The spectrum for ferric citrate was largely similar to that for citrate at pH above 3.5 with the symmetric, asymmetric, and carbonyl stretches present at similar locations (Figure
2.3). Differences exist, however. For example, the decreased intensity of the carbonyl stretch relative to the C-Osym and C-Oasym stretches indicated the carboxylic acid groups deprotonated at a lower pH in the presence of iron. The feature at 1280 cm-1 associated with the C-O-H hydroxyl bending motion was also clearly reduced in the presence of bound ferric iron. Finally, the two peaks bracketing the 1280 cm-1 peak, which were related to C-H rocking, were more clearly defined peaks as opposed to the shoulders observed in the citric acid spectra. The changes in the spectra at ~1280 cm-1 presumably reflect the hydroxyl group deprotonating and taking an active role in binding ferric iron in the complex. The deprotonation of the hydroxyl group was consistent with previously reported structures of ferric citrate crystals precipitated from circumneutral pH solutions where citrate was present in excess of iron (Matzapetakis et al., 1998; Pierre and Gautier-
Luneau, 2000). Similar ferric structures have also been observed in solution over a range of pH values with the hydroxyl group deprotonating more as the pH increases (Vukosav et al., 2012).
2.3.2.2 Adsorbed Citrate spectra
Similar to spectra for aqueous citrate and ferric citrate, those for citrate adsorbed on the
LSA hematite were dominated by carbonyl, C-O asymmetric and C-O symmetric stretches across a range of pH values (Figure 2.4). However, there were several subtle 38 differences between the aqueous and adsorbed citrate that provide insight into the structure of the adsorbed complex. To begin, although the dominant peaks show little shift in position they do exhibit an increase in peak width. The similarity in the position of the major peaks between the reference solution phase spectra and those for adsorbed citrate suggests citrate was electrostatically bound as an outer sphere complex (Dobson and McQuillan, 1999; Roddick-Lanzilotta and McQuillan, 2000). This was consistent with the reduction in citrate adsorption at low pH with increasing ionic strength (Figure
2.2). According to Roddick-Lanzilotta and McQuillan (Roddick-Lanzilotta and
McQuillan, 2000), broadening of IR peaks, with no change in position, results from the formation of complexes where the carboxyl groups’ conformation remain similar to that formed in the aqueous state, such as outer sphere complexes or a bridging multi-dentate inner-sphere complex. The presence of the symmetric and asymmetric C-O stretches in the pH 2.5 adsorbed citrate spectrum indicates citrate was partially deprotonated at the surface since these features were not present in the corresponding aqueous-phase spectrum (Figure 2.1). This was consistent with the expected shift in the interfacial pH towards the pHpzc (Johnson et al., 2004). Using the diffuse layer model, Dzombak and
Morel (Dzombak and Morel, 1990) demonstrated the pH of a 100 mM 1:1 electrolyte solution will shift from 2.5 in the bulk to around 5 near the surface. A similar shift in our system would take citrate from a fully protonated state in the bulk solution to a mixture of roughly 67% doubly protonated and 25% singly protonated citrate near the surface.
Evidence for the presence of an additional complex was manifest in the spectra when the concentration of citrate was varied (Figure 2.5). At a total citrate concentration of 62.5
µM at pH 3, nearly 100% of the total citrate was adsorbed. The resulting spectrum shows 39 -1 no carbonyl peak at 1720 cm , however the νasC-O and νsC-O were present and had peak heights of roughly 30% and 50% those at the maximum citrate concentration of 500 µM.
As the citrate concentration was increased, the carbonyl peak gradually appears suggesting the presence of an additional protonated complex, likely outer-sphere. The absence of a carbonyl peak at low citrate concentrations likely reflects at low surface coverages that citrate preferentially adsorbs in this deprotonated manner to a limited number of sites. Increasing surface coverage consequently results in the formation of additional outer-sphere complexes that were protonated. Evidence of such a preferential adsorption mode of citrate at specific hematite surface sites was supported by the observation of Cornell and Schwertmann (Cornell and Schwertmann, 2003b) that adding citrate to an Fe(III) solution during hematite synthesis causes the hematite to grow along the (001) face due to the specific adsorption of citrate to the (110) and (104) surfaces.
The (001) face consists of only doubly coordinated hydroxyl groups and it is neutrally charged at environmentally relevant pH values (Barron and Torrent, 1996). The (110) and (104) faces consist of singly, doubly and triply coordinated hydroxyl groups, which appear to be preferential for citrate adsorption (Barron and Torrent, 1996). The outer sphere complex was evident at elevated pH values as well based on the peak broadening and lack of shift in peak position compared to the aqueous citrate spectra (see Figure 2.4).
At pH 6, there was also evidence of an inner-sphere structure in the feature at 1280 cm-1 which arises from the C-O-H bending motion. This feature was nearly absent at 62.5 µM citrate, indicating that either the hydroxyl deprotonates and takes an active role in the adsorption process as seems likely when complexing ferric iron (Matzapetakis et al.,
1998; Pierre and Gautier-Luneau, 2000) (see Figure 2.3) or due to its proximity to the 40 surface it was constrained and prevented from moving freely. Both options require direct binding to the surface. As the concentration was increased, this feature becomes more obvious; however, at maximum surface coverage it was not as prominent as it was in the aqueous citrate spectra (see Figure 2.3).
The high surface area hematite spectra (Figure 2.4) show a similar pattern as the LSA hematite. The biggest difference between the two was in the carbonyl stretching region at low pH where the HSA spectra had a much smaller carbonyl stretch. The vasC-O and vsC-O features for both HSA and LSA were comparable. The carbonyl stretch in the HSA spectra was nearly gone by pH 3.46 whereas in the LSA spectra it persists until pH 5.5.
As the variable citrate concentration LSA spectra showed, citrate has a preferential fully deprotonated inner sphere adsorption mode that dominates at lower citrate concentrations before citrate begins adsorbing in an outer-sphere manner at higher concentrations. The subdued carbonyl stretch relative to that in the LSA spectra suggest that the HSA hematite had a higher percentage of sites that favor inner-sphere complex formation.
Madden et al. (Madden et al., 2006) report that high surface area hematite particles prepared using similar methods show an increased affinity with copper(II) ions when compared to a lower surface area hematite. This they attribute to an increase in the number of surface sites on the higher surface area hematite that suit the distorted octahedron binding environment that Cu(II) prefers. The differences in citrate adsorption observed for LSA and HSA hematite could reflect similar differences in available surface sites.
41 2.3.3 Computational Modeling
-3 The computed spectra for citric acid (H3Cit) and citrate (Cit ) (Figure 2.3) did not require a scaling factor in order to match the experimental symmetric and asymmetric peak locations, therefore no scaling factor was applied to either the aqueous or the adsorbed citrate results. Explicit water molecules were required, however, as their inclusion resulted in the computed spectra peak position and intensity being closer to those in the aqueous citrate experimental spectra (comparison shown in Figure A.3 in Appendix A).
The major area of improvement was observed for the C-O-H bending of the hydroxyl group, which in the absence of explicit water molecules had a calculated intensity equal to that for the C-O symmetric stretch. Experimentally, however, the C-O-H bending feature appears as a shoulder to the symmetric C-O stretch at ~1443 cm-1 in the experimental spectra (Figure 2.3).
Proposed structures for adsorbed citrate were informed by the batch adsorption and spectroscopic data. The FTIR results suggest outer-sphere or binuclear bidentate inner- sphere complexation modes and thus these were the initial, though not sole, focus of the computational simulations. Both singly protonated and fully deprotonated conformations were simulated as the appearance of the carbonyl stretch in the pH 3 adsorbed citrate spectra at ~1720 cm-1 with increasing surface coverage (Figure 2.5) suggested both could exist. The peak locations of the experimental and theoretical spectra are summarized in
Table 2.3 and plotted in comparison to one another in Figure A.4. Select theoretical spectra are shown in Figure 2.6 (structures shown in Figure A.2) using a Lorentzian distribution with the full width at half height of 20 cm-1 for all peaks (Kawiecki et al.,
42 1988). As the signal from the water and hematite were subtracted from the experimental
FTIR spectra by the subtraction of standards, the features arising primarily from the iron cluster and explicit water molecules were excluded from the computed spectra using
Gaussian 09.
The mononuclear bidentate complex (see CT-MN in Figure A.2) structure had citrate bound to a single iron atom by the central carboxyl (the carboxyl group attached to the central carbon on the citrate backbone, adjacent to the hydroxyl group) and one terminal carboxyl group (one of the two on either end of the molecule). This arrangement does not match the experimental data well. The peaks representing νasC-O and νsC-O were further apart than in the experimental data and the νasC-O and νsC-O were split into several distinct peaks. Based on this, the likelihood that mononuclear complexes formed at the hematite surface appears low.
The location of the carbonyl stretch, which was only present in the experimental spectra below pH 5.5, was not correct in any of the inner-sphere theoretical spectra. The bidentate structure involving the deprotonated hydroxyl group (CH-BN in Figure A.2) produced a theoretical carbonyl stretch at 1665 cm-1, well below the LSA experimental value of 1720 cm-1, while the central and terminal carboxyl bound model (CT-BN in
Figure A.2) errs the other way at 1749 cm-1. The outer-sphere structure (OS-BN in
Figure A.2) produced a carbonyl stretch that was much closer at 1726 cm-1. This, along with the influence of ionic strength on adsorption at low pH (see Figure 2.1) indicates the protonated citrate complex was likely an outer-sphere complex.
43 The fully deprotonated bidentate model that included a deprotonated hydroxyl group interacting with the iron cluster and the outer-sphere models were the atomic configurations that most accurately matched the experimentally determined symmetric and asymmetric C-O stretches. The νsC-O of the protonated outer-sphere complex was at
1411 cm-1 (1410 cm-1 for the deprotonated OS complex) and for the inner-sphere complex with the deprotonated hydroxyl group it was at 1398 cm-1. There was a minor shift in the position of νsC-O of adsorbed citrate in the experimental LSA spectra from
1403 to 1394 cm-1 as pH increased, likely related to the sequential deprotonation of the citrate. This shift was not seen in the HSA hematite spectra where the vsC-O for adsorbed citrate remained at ~1398 cm-1 over the entire range of pH tested (see Figure 2.4), indicating that citrate adsorbed to HSA hematite was fully deprotonated at a lower pH and possibly indicating inner-sphere binding was more dominant than outer-sphere binding. This shift in peak location was seen as an additional shoulder in the data presented by Lindegren et al. (Lindegren et al., 2009) for citrate on goethite and it was attributed to conformational changes in the adsorbed citrate.
2.3.4 Surface complexation modeling
Surface complexation modeling was performed to further test the structures determined from the match of the theoretical spectra to the experimental spectra. The triple layer model was used with published constants for surface protolysis and electrolyte binding for both sizes of hematite (Hwang and Lenhart, 2008) with citric acid solution chemistry taken from Stumm and Morgan (Stumm and Morgan, 1996) and Silva et. al. (Silva et al.,
2 2009) (see Table A.1). The site density, Ns, was set at 2.3 sites/nm as fitting the data to 44 2 2 a Langmuir isotherm yields a Γmax of 1.83 µmol/m or 2.2 sites/nm , very close to the recommended and commonly used 2.3 sites/nm2 (Davis and Kent, 1990). The stoichiometry shown in eqs. 5-7 of Table 2.3 represent the best fitting outer- and inner- sphere species being considered. Fits to these reactions were conducted for a citrate concentration of 1000 µM on the LSA hematite, which produces a saturated condition on the hematite surface. These constants were then fixed and the model was tested against different solution condition data sets. The binuclear stoichiometry was used as it was the most likely arrangement resulting from the FTIR analysis and molecular modeling and was also consistent with studies of citrate on goethite (Lindegren et al., 2009) and corundum (Hidber et al., 1996). Two outer-sphere complexes were tested, a singly protonated species and a fully deprotonated species, both bound to two fully protonated
FeOH2 sites (See Figure A.5 for simplified cartoon). The entire charge for the outer- sphere citrate complex was placed in the beta layer. Additional simulations that distributed the charge to different planes did not result in marked improvements in model fits. The inner-sphere complex included a completely deprotonated citrate molecule, with a charge of -4 (See Figure A.5 for simplified cartoon). The reaction was modeled as a ligand exchange reaction with all of the complex charge directly at the surface. As was done for the outer-sphere species, simulations placing the charge of the inner-sphere species at different planes did not improve the model.
Fits were performed using just the outer-sphere complexes or with a combination of the outer- and inner-sphere surface complexes. The WSOS/DF for the outer-sphere only model on the LSA hematite was 1.92. Including the inner-sphere complex only reduced it
45 slightly to a value of 1.33, but were more consistent with the spectroscopic and molecular modeling results. Simulations with this dual-mode model indicated the singly protonated outer-sphere complex was dominant, accounting for approximately 90% of the adsorption at low pH (see Figure 2.7a). This continued as the pH was raised until the deprotonated outer-sphere complex became dominant at a pH of 8.5. This description of the system appears to vary slightly from the adsorbed citrate FTIR data in that the carbonyl peak at 1720 cm-1, the indicator of protonated carboxyl group, disappears above a pH of 5.5 (see Figure 2.4). At this pH, the SCM model indicated the singly protonated outer-sphere citrate represented 90% of the total adsorbed citrate (Figure 2.7a). Although this pH was below the third pKa for citric acid, where over half of the citrate was present in the singly protonated state, the carbonyl peak was only slightly visible in the aqueous citrate IR spectrum (Figure 2.3). Thus, while the SCM probably overestimates the singly protonated OS species, it was not a completely unreasonable fit as the protonated species can exist in significant concentrations with only a small carbonyl peak visible in the spectrum. There is precedence for a slight discrimination between the FTIR analysis and
SCM fitting. Lackovic et al. (Lackovic et al., 2003) also paired FTIR and SCM in investigating citrate on goethite and also had a singly protonated complex comprising the majority of the adsorbed species and existing well beyond the final pKa of citrate. Hwang et al. (Hwang et al., 2007) used FTIR to determine phthalate forms two inner sphere and one outer-sphere complexes on hematite, all deprotonated. The SCM fits following the
FTIR analysis were accomplished using only one inner-sphere and one outer-sphere complexes, all deprotonated (Hwang and Lenhart, 2009).
46 When these three adsorption reactions and related equilibrium constants were applied to data from a solution with 278 µM citrate, the quality of fit decreased (WSOS/DV = 37.4).
At this citrate concentration, which results in slightly less than full surface coverage, the inclusion of the inner-sphere complex becomes more important and at low pH it accounted for 25% of the total adsorption (Figure 2.7b). This was in agreement with the idea that at low pH the inner-sphere complex forms preferentially to a limited number of higher-affinity sites and as the citrate concentration increased and these sites were saturated the formation of outer-sphere complexes became important. At a lower electrolyte concentration (2 mM NaCl), the model prediction to the data produced a
WSOS/DF value of 59.2 as the model fit the data well at low pH, but failed to capture the adsorption accurately at high pH values (Figure 2.7c). This may be due to the presence of an additional inner-sphere complex that was not accounted for in the model. The inclusion of an additional inner-sphere complex would be consistent with the lack of change in the amount of citrate adsorbed at high pH as the ionic strength was varied (see
Figure 2.2). It has been suggested that citrate forms an inner-sphere complex on goethite
(Lindegren et al., 2009) and corundum (Hidber et al., 1996) in the slightly basic pH range and it was possible that is also the case on hematite, however, including an extra complex to our SCM model did not produce improved fits. In keeping with the proposed dominance of the outer-sphere complexes, the model predicts that lowering the ionic strength will cause an increase in the abundance of the outer-sphere complex, especially at low pH where the OS complex is increased from 0.55 µM at 100mM NaCl to 1.29 µM at 2 mM NaCl.
47 Sverjensky, (Sverjensky, 2003, 2005) proposed the following method for adjusting equilibrium constants to account for variations in surface area that result when varying particle size:
Log Kθ = Log K0 NSAS / N#A#
Where Ns is the site density for the sample, As is the specific surface area of the sample and Kθ is the equilibrium constant for the sample. N#, A#, and K0 are the same values for the hypothetical standard state. In addition to using a different equilibrium constant for the HSA hematite, a new inner-layer capacitance, C1, also provided by Hwang and
Lenhart, (Hwang and Lenhart, 2008) must be used as C1 has an inversely proportional relationship with specific surface area (Hwang and Lenhart, 2008). Using the LSA
2 hematite with N# set to 2.3 sites/ nm to determine the equilibrium constants for the HSA sample and the parameter values from Hwang and Lenhart (Hwang and Lenhart, 2008) summarized in Table 2.1, resulted in a model predicted WSOS/DF of 15.3 for the dual- mode model applied to the HSA hematite with 1000 µM citrate (Figure 2.7d). As expected for the smaller hematite, the lower C1 increased the importance of the inner- sphere complex as a smaller C1 physically implies a greater distance between the surface and the beta plane (Sverjensky, 2001). The increase in inner-sphere complexation as the particle size was reduced was consistent with both the spectroscopic data presented here and previously reported dual-mode model fits of phthalic acid adsorption on hematite
(Hwang and Lenhart, 2009). However, it contradicts the model proposed by Lackovik et al. (Lackovic et al., 2003) for citrate on goethite using an extended constant capacitance model. In their model, a singly protonated mononuclear monodentate inner-sphere 48 complex dominates, accounting for nearly all adsorption below pH 7.7. Above pH 7.7, a deprotonated outer-sphere complex also arises. Although the model of Lackovik et al.
(Lackovic et al., 2003) is very different from ours, it also failed to accurately match adsorption data at high pH values. This suggests an additional species that was not accounted for in our modeling effort, or those previously conducted, might exist. Thus, additional effort is needed to determine the nature of the additional adsorption mode at high pH.
2.4 Conclusion
Explicitly defining details of the adsorption of organic molecules like citrate on iron oxides is not a straightforward problem. For goethite, many methods have been used to study citrate adsorption, but as of yet a consensus on the exact adsorption mechanism does not exist. The four methods of determining the structure of citrate on the hematite surface used in this study all point to outer-sphere complexes dominating the adsorption of citrate on hematite surfaces across all pH values. These complexes differ on the basis of protonation state, with the outer-spherically bound citrate molecule changing from singly protonated to a fully deprotonated complex at a pH of approximately 5.5 - 6.5.
Supplementing these species was a fully deprotonated binuclear bidentate inner-sphere complex. The inner-sphere complex forms at lower citrate concentrations while the outer-sphere complex doesn’t form until citrate was present in higher concentrations.
This reflects the limited amount of inner-sphere sites relative to those available for the electrostatically held citrate. The inner-sphere mode appears to be slightly more significant in the HSA hematite, indicating the proportion of the sites to which citrate 49 adsorbs directly does not change in exact proportion with the overall hematite surface area. Testing a broader range of hematite particle sizes and morphologies is necessary to arrive at a definitive conclusion on the impact of particle dependent reactivity.
Acknowledgments:
We would like to thank the National Science Foundation under Grant No. 0954991 for providing funding for this research. We also acknowledge the Ohio Supercomputer
Center for use of their computational resources.
50
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57
1 2
a) HSA ) b) HSA 2 LSA LSA 0.8 1.5 mol/m µ 0.6 1 0.4
0.5 Fraction Adsorbed 0.2 Surface Coverage ( 0 0 2 4 6 8 10 2 4 6 8 10 pH pH Figure 2.1. Adsorption envelope for citrate (1000 µM) adsorbed onto LSA and HSA hematite at saturated conditions presented as a) fraction adsorbed and b) surface coverage.
58 2 [NaCl] = 0 mM 1.8 [NaCl] = 100 mM )
2 [NaCl] = 500 mM 1.6
1.4 mol / m µ 1.2
1
0.8
0.6
0.4 Surface Coverage ( 0.2
0 2 4 6 8 10 pH Figure 2.2. Surface coverage of adsorbed citrate on the low surface area hematite as a function of pH at 3 different NaCl concentrations.
59 Ferric Citrate
pH=8.5
pH=5.5
pH=2.5
Citrate
Citric acid
1800 1700 1600 1500 1400 1300 1200 −1 Wavenumber (cm ) Figure 2.3. Reference spectra used for comparison to adsorbed citrate spectra. The ferric citrate spectra at 3 pH values represent models for inner-sphere complexation. Additional experimental spectra (top, black) are presented for fully protonated citric acid (pH 2.5) and fully deprotonated citrate (pH 9.5). Corresponding theoretical spectra are also presented (blue, bottom) where the vertical lines (red) represent the individual absorbance frequencies from the DFT calculation.
60 C=O as sym C−OH
Citrate
LSA pH 8.51
HSA pH 8.59
LSA pH 5.84
HSA pH 5.80
LSA pH 2.51
HSA pH 2.43
Citric acid
1800 1700 1600 1500 1400 1300 1200 −1 Wavenumber (cm ) Figure 2.4. Spectra for adsorbed citrate on LSA and HSA hematite at given pH values (1 mM citrate).
61 a) pH = 3.0
b) pH = 6.5
1800 1700 1600 1500 1400 1300 1200 −1 Wavenumber (cm ) Figure 2.5. FTIR spectra of citrate adsorbed on LSA hematite at (a) pH 3.0 and (b) 6.5. The citrate concentrations for the spectra from bottom to top for both pH values were 62.5, 125, 250, 500 µM.
62 OS−BN
OS−BN−DP
CT−MN
CT−MN−DP
CT−BN
CT−BN−DP
CH−BN
CH−BN−DP
1800 1700 1600 1500 1400 1300 1200 −1 Wavenumber (cm ) Figure 2.6. Selected theoretical infrared spectra for adsorbed citrate structures visualized with a Lorentzian distribution with a 20 cm-1 FWHM. The short red vertical lines represent the absorption energies. The dashed vertical lines reflect the peak locations for the experimental results of the low surface area hematite at mildly acidic pH. (OS = outersphere, MN=mononuclear, BN = Binuclear, DP = deprotonated, CT=central and terminal carboxyl bonding, CH=Central and hydroxyl bonding.) The optimized structures are shown in supporting information (Figure A.2)
63
A) B)
) LSA Hematite LSA Hematite 2 [Cit] = 1000 M [Cit] = 278 M 1.5 OS µ 1.5 µ P [NaCl] = 100mM [NaCl] = 100mM
mol / m OS P µ 1 1
0.5 0.5 OS OS DP IS DP IS Surface Coverage (
0 0 2 4 6 8 10 2 4 6 8 10
C) D)
) LSA Hematite HSA Hematite 2 1.5 [Cit] = 278µM 1.5 [Cit] = 1000µM [NaCl] = 2mM [NaCl] = 100mM
mol / m OS P µ OS P 1 1
0.5 0.5 IS OS IS OS DP DP Surface Coverage (
0 0 2 4 6 8 10 2 4 6 8 10 pH pH Figure 2.7. Results of surface complexation modeling using one inner-sphere (IS), one singly protonated outer-sphere (OSP) and one deprotonated outer-sphere (OSDP) complex as described in eq 5-7 of table 2.1. Individual species are shown with dashed lines and the total adsorption with a solid line. Equilibrium constants were determined from the data in plot A and applied to data collected at other solution conditions (B-D).
64
Table 2.1. Constants used for triple layer model fitting. All values taken from Hwang and Lenhart(Hwang and Lenhart, 2008) aside from the results of the best model fit of 1000 µM Citrate to the LSA hematite (eq 5-7). Values given as logs of equilibrium constants.
N s pH C (F/m2) C (F/m2) SSA (m2/g) (sites/nm2) pzc 1 2 LSA 2.3 9.05 0.92 0.2 35 HSA 2.3 9.17 0.59 0.2 99 eq Log K LSA HSA + - + 1 >FeOH + H => >FeO + H Ka1 6.42 6.20 + + 2 >FeOH + H => >FeOH2 Ka2 -11.68 -11.92 + - + + 3 >FeOH + Na => >FeO --- Na + H Kcation -8.58 -8.32 - + + - 4 >FeOH + Cl + H => >FeOH2 --- Cl Kanion 9.47 9.90 + 2- 5 (>FeOH)2 + H3Cit => (>FeOH2 )2 --- HCit KOS1 20.23 20.68 + 3- + 6 (>FeOH)2 + H3Cit=> (>FeOH2 )2 --- Cit + H KOS2 13.70 14.15 2- + 7 (>FeOH)2 + H3Cit => >Fe2H-1Cit + 2H2O + 2H KIS -2.23 -1.78
65
Table 2.2. DFT assignments of simulated peaks for citrate and citric acid
Citric Acid Citrate3- wavenumber assignment wavenumber assignment 1745.97, 1753.95, 1768.01 ν C=O 1572.50, 1591.06, 1596.98 νC-O(asym) 1420.56, 1444.98 δCH2, δCOH 1449.69 δCOH 1175.51, 1220.23, 1245.76 δCH2, δCOH 1402.30, 1407.60, 1425.07 νC-O(sym) 1290.40, 1331.70 δCH2 1311.75 δCH2, δCOH
66
Table 2.3. Experimental and theoretical symmetric and asymmetric stretch peak locations. The protonated mononuclear tridentate structure did not exhibit either a symmetric or asymmetric stretch. Structures for theoretical complexes are show in Figure A.2.
System pH νC=O νC-Oas νC-Osym Citrate 1.5 1724.9 1398.7
3.5 1723.2 1578.8 1397.5
5.6 1725.9 1569.7 1390.3
8.7 1567.3 1390.1 FeCit 3.5 1590 1385.5
5.5 1577.2 1389.8
8.5 1567.3 1388.9 10 nm Hem 2.4 1720.4 1578.7 1398.3 5.8 1573.3 1397.3
9.5 1579.5 1397.2 50 nm Hem 2.5 1715.2 1579.1 1403.4 5.8 1714.3 1576.8 1397.3
9.5 1577.6 1394.9
Theoretical IS-Mononuclear 1754.4 1641.7 1370.9
IS Mononuclear Deprotonated 1601.9 1302.7
IS Bidentate 1748.5 1590.8 1414.4
IS Bidentate Deprotonated 1597.3 1436.9
IS Bidentate Hydroxyl 1664.4 1581.1 1371.4
IS Bidentate Hydroxyl DP 1583.9 1398.3
OS 1726.2 1606.8 1410.0
OS Deprotonated 1576.1 1411.7
67
Chapter 3: X-Ray Analysis Of Lead Adsorbed On The Hematite (001), (012), And (110)
Surface
Abstract:
Two synchrotron based techniques, Extended X-ray adsorption fine structure (EXAFS) and X-ray Reflectivity (XR) with resonant anomalous X-ray reflectivity (RAXR) were used to determine the adsorption of lead on the hematite surface. EXAFS was performed on two different sizes of hematite while XR/ RAXR is performed on single crystal surfaces expressing only one crystal face. The EXAFS showed that the lead adsorbs in a bidentate inner-sphere manner in both edge sharing and corner sharing arrangements on the FeO6 octahedra. Three hematite faces were tested, (001), (012), and (110), which are considered the most common faces. The reflectivity experiments confirmed the inner sphere adsorption modes and revealed additional outers-sphere adsorption modes that were not seen in the EXAFS. The (001) surface was the least reactive surface and did not change with pH. The (012) surface was the most variable with pH and had the highest surface coverage at pH 6. At pH 4, the (110) had the highest surface coverage. The (012) and (110) surfaces have singly and triply coordinated oxygen atoms which are likely the source of the increased reactivity.
68 3.1 Introduction
Lead is a highly toxic element that negatively influences nearly all bodily systems (Casas and Sordo, 2006) and due to decades of use as an additive to gasoline and paint it is among the most widely dispersed metal pollutants in the environment being found from remote American forests (Kaste et al., 2006) to the ice sheets in Greenland and
Antarctica(Boutron et al., 1994), though most heavily in urban areas (Mielke et al.,
2011). Lead has a strong affinity for metal oxide minerals and as such it is not expected to remain in solution in soil systems (Hassellov and von der Kammer, 2008; O'Reilly and
Hochella Jr, 2003). Adsorption reactions are therefore important to the fate and transport of lead through the environment. Lead has been found associated with iron oxides in aquatic (Taillefert et al., 2000) and soil systems (Kaste et al., 2005; Kaste et al., 2006;
Schroth et al., 2008) allowing for the possibility that lead may be adsorbing onto nano- scale iron oxide particles which remain mobile (Hassellov and von der Kammer, 2008;
Lofts and Tipping, 2000).
Hematite, α-Fe2O3, is a naturally occurring iron oxide commonly found in soils in both the nano and macro scale (Cornell and Schwertmann, 2003). Hematite is one of the more thermodynamically stable iron oxides and it is the end transition form of other less stable iron oxides (Cornell and Schwertmann, 2003). Particles of hematite may exhibit a variety of crystal morphologies, from well defined platy crystals to more spherical shapes, depending on the conditions under which they were formed (Cornell and Schwertmann,
2003). The different particle morphologies result in a different distribution of crystalline faces each of which may exhibit unique charging and adsorbent characteristics (Venema
69 et al., 1998). Knowledge of the crystal faces for hematite is not as robust as for other iron oxides such as goethite. However it is generally agreed that the most common crystal faces are (001), (012) and (110) (Catalano et al., 2009; Mackrodt et al., 1987).
Lead adsorption on hematite particles has been studied using a number of methods both experimental (Bargar et al., 1997; Bargar et al., 2004; Lenhart et al., 2001) and theoretical(Mason et al., 2009). From macroscopic batch adsorption studies, we know that the pH dependent adsorption edge of lead on hematite increases sharply between approximately pH 4 and 6 depending on the electrolytes present in solution and solid:solution ratio (Bargar et al., 1998; McKenzie, 1980). McKenzie (McKenzie, 1980) reported the maximum surface coverage of lead on hematite at pH 5 was 2.7 µmol/m2 and that the adsorbed lead was not easily extractable with acetic acid, as it was from the goethite surface indicating that the lead was held by a strong inner-sphere complex.
Spectroscopy provides direct information on adsorbed structures and Bargar et al.(Bargar et al., 1997) utilized extended X-ray absorption fine structure spectroscopy (EXAFS) to determine that lead adsorbs as a mononuclear bidentate inner-sphere complex on the edge of the iron oxide octahedron at pH 6-8 as evidenced by the lead iron distance of 3.3 Å.
Also using EXAFS, Lenhart et al. (Lenhart et al., 2001) corroborated the edge sharing structure and found an additional binuclear bidentate corner sharing arrangement of lead on the hematite surface at pH 6 with a lead – iron distance of 3.8 Å. While EXAFS provides details of the local coordination environment around the lead, it can be difficult to determine exact binding mechanisms as it provides an average of the coordination environment in the sample. This means that details regarding binding at specific
70 crystalline faces, contributions from minor coordination modes, etc. are lost during sample analyses and data fitting (Nelson and Miller, 2012).
Crystals can either be cut to a specific surface or grown on a substrate causing a single surface to be exposed. The exact makeup of the surface is then known allowing for studies to more precisely elucidate adsorption mechanisms. There have been several attempts made at investigating lead adsorption on these single crystal surfaces (Bargar et al., 2004; Mason et al., 2009). Bargar et al.(Bargar et al., 2004) used Grazing Incidence
EXAFS and X-ray photoelectron spectroscopy to investigate hematite (001) and (012) surfaces, finding high lead surface coverage in the form of an oligomeric complex on both. Mason et al.(Mason et al., 2009) used density functional theory (DFT) computation methods to investigate the coordination of lead on the (001) surface of hematite and the isostructural aluminum oxide, corundum, finding that the lead is bound four times stronger to hematite. Lead coordination was in a tridentate fashion with Pb-O distances slightly shorter than the EXAFS results (Bargar et al., 1997; Lenhart et al., 2001). The exact distance in the DFT models depended on how the protons were distributed on the hematite surface at the beginning of the simulation.
The particle based (EXAFS) and the single crystal X-ray reflectivity offer different strengths and weaknesses. EXAFS is highly precise, however it is element specific. The results of an EXAFS analysis yield a description of the environment immediately around the target atom. In the case of lead, the EXAFS is only able to probe at most 5 Å from the lead atom. In a heterogeneous system, all of the local coordinating environments surrounding lead are averaged in the result. X-ray reflectivity (XR) is a surface specific
71 technique. The information gleaned from reflectivity is the electron density as a function of distance from the surface (Fenter, 2002). XR can be combined with resonant anomalous X-ray reflectivity (RAXR) which gives the electron density of a specific element as a function of distance from the surface (Lee et al., 2011). As a result, reflectivity allows us to easily see the difference between inner sphere and outer-sphere adsorption process. This method requires that the surface be cut or cleaved to a specific plane. Knowing exactly what the adsorbent surface is adds clarity to the adsorption process at the expense of direct environmental relevance.
In this paper, the two different synchrotron-based X-ray techniques were used to better elucidate the coordination environment of lead at the hematite surface. EXAFS was performed at pH 6, confirming the findings of previous studies that lead adsorbs in an inner-sphere manner to the hematite surface forming both corner sharing and edge sharing bidentate complexes on the FeO6 octahedra. XR / RAXR was performed on the
(001), (012), and (110) surfaces at pH 4 and 6. The total surface coverage varied between the surface following (012) > (110) > (001) at pH 6 and (110) > (012) > (001) at pH 4. The lead was held to the surface by a combination of inner- and outer-sphere complexes. By combining both methods we can get a more complete picture of lead adsorption on hematite.
3.2 Experimental
3.2.1 Hematite
Both hematite particles and single crystals were used in this series of experiments. Two sizes of hematite particles were synthesized, nominally 50 nm and 10 nm, as described in 72 detail in Chapter 2 and in the supporting information. Crystals, oriented, cut and polished to expose the (001) surface were procured from a natural hematite crystal and provided courtesy of Dr. Glenn Waychunas (Lawrence Berkeley National Laboratory). Crystals similarly prepared to display the (012) and (110) surfaces were purchased from
SurfaceNet GMBH. The crystal surfaces were cleaned prior to each experiment following methods of Catalano et al.(Catalano et al., 2007a; Catalano et al., 2008) by placing each crystal in a methanol bath and sonicating for 5 minutes, followed by 5 minutes in an acetone bath. This was repeated 5 times. Following this, the crystals were placed in alternating 10-3 M NaOH and HCl baths and sonicated. Crystals were rinsed with DI water (18.2 MΩ from Millipore system) between baths. Following the final acid bath, the crystals were completely rinsed with DI water and annealed at 450oC for 4 hours and slowly cooled (Catalano et al., 2007a; Catalano et al., 2008).
3.2.2 EXAFS
Samples were prepared for EXAFS analysis by mixing the hematite nanoparticles with lead and a sodium perchlorate sufficient to bring the final concentration of lead to 0.7 mM for the LSA and 1 mM for the HSA and sodium perchlorate to 0.1 M in polycarbonate centrifuge tubes. The solutions were pH adjusted using NaOH or perchloric acid. All solutions were CO2 free and all transfers were performed under a humidified nitrogen atmosphere. Samples were equilibrated on an end-over–end rotator in the dark for 24 hours when the pH was checked and adjusted if necessary and returned to the rotator. At 48 hours, the samples were removed from the rotator and the final pH was determined before the samples were centrifuged to separate the solids from the
73 supernatant. The particle wet pastes were immediately removed and mounted into PTFE holders that were sealed with Kapton tape. The mounted samples were kept in a cool damp nitrogen atmosphere until they were measured. The lead concentration of the supernatant was analyzed with an inductively coupled plasma – atomic emission spectrometer (Varian Vista AX CCD-Simultaneous ICP-AES) to determine the lead surface coverage.
An EXAFS spectrum is created by measuring the absorbance of the x-ray beam as the energy is varied across the binding energy of a specific electron, in this case the lead LIII electron with a binding energy of 13035 eV. The spectrum is created when the energy is equal to or above the edge energy causing the target electron to be ejected from the atom, becoming a photoelectron and creating a core hole. The photoelectron is ejected at the same difference in energy between the incident x-ray and the binding energy and it is free to interact with neighboring atoms. Depending on the energy of the photoelectron it can scatter from neighboring atoms back into the abandoned core hole thus affecting the absorption coefficient. Because 1) there is a very limited amount of time, on the order of femtoseconds, before an electron from a higher energy level falls down into the core hole and 2) the probability of the photoelectron reaching an atom and returning to the original atom decreases with distance, EXAFS provides information only very close to the target atom (Newville, 2014). Lead, being a very deformable atom will only yield information at most 5 Å from the lead atom with very high quality data.
Pb LIII EXAFS spectra were collected in fluorescence mode at beamline 20-BM-B at the
Advanced Photon Source at Argonne National Laboratory. A Si 111 monochromator
74 was used to select the x-ray energy and a 1 x 6 mm unfocused incident beam limited by beam defining slits was applied to the samples. The x-ray beam was detuned 15% to eliminate harmonics and a Canberra 13-element germanium detector was used to measure the samples’ fluorescence. Several layers of aluminum foil were placed between the sample and detector to eliminate the Fe Ka fluorescence. The detector was set to measure the fluorescence from the L3-M5 transition at 10551 eV. At least 10 scans of all samples were taken and merged together to improve the signal to noise ratio. Data analysis was performed using Athena and Artemis in the Demeter package (Ravel and Newville,
2 2005). To fit the data, a fixed amplitude reduction factor, So , of 0.8425 was used for all elements (Lenhart et al., 2001). The Debye-Waller function was set to 0.01 for all oxygen and iron atoms (Bargar et al., 1997; Lenhart et al., 2001). The electron binding energy shift, Eo, was determined by allowing it to vary in the fit of the first shell oxygen atom and then setting the Eo at the determined value for the rest of the fit (Lenhart et al.,
2001). Input models to FEFF were created by building a structure consisting of a lead atom on an iron oxide octahedra.
3.2.3 X-ray Reflectivity
All reflectivity data was collected in situ at Beamlines 33ID and 6ID at the Advanced
Photon Source. Reflectivity experiments were designed to collect two complimentary sets of data, nonresonant X-ray reflectivity (XR) analysis to determine the overall electron density and resonant anomalous X-ray reflectivity (RAXR) to determine the lead specific electron density. The two data sets were collected in series without any changes to the setup aside from flushing the crystal surface with fresh Pb solution. The samples
75 were prepared by placing clean crystals into a freshly prepared 0.1 mM Pb solution with a
0.1 M sodium perchlorate as a background electrolyte, both free of CO2. Each surface was tested at pH 4 and pH 6 as the lead adsorption edge on hematite is located within this pH range (McKenzie, 1980). After equilibration, the crystals were placed in a flow through cell with a Kapton film window (Bellucci et al., 2015). The solution used during the crystal surface equilibration was also utilized to flush the cell to remove any air bubbles. The solution was allowed to drain to minimize the amount of solution on the crystal surface to reduce the attenuation of the x-ray beam. The holder was placed in the diffractometer and the crystal aligned so that both miscut reflections were visible and centered.
An in-depth description of the XR and RAXR measurement and data analysis can be found in Appendix C. Briefly, XR data was collected by varying the electron momentum transfer (q = 2πL/d where L is the Bragg index and d is the length of the unit cell) while keeping the energy constant at 12 keV, an energy not near absorption edges of any of the elements involved. System stability was monitored in the XR by either returning to a specific q value several times during the course of data collection (q = 1.23 and 2.46 on the (001) surface only for data collected at 33-ID) or by measuring every other data point first from high to low q, then back from low to high q to fill in the skipped data points
(for data collected at 6-ID).
The RAXR scans were collected by varying the photon energy across the LIII edge of Pb at 13.035 keV while maintaining a constant q. This was repeated for multiple values of q. The spectrum at a specific low q value (q ~ 0.6 Å-1) was measured repeatedly
76 throughout the collection of spectra. If the spectrum changed, the sample was flushed with fresh solution and the sample moved so that a new spot was illuminated by the X- ray beam. The stability of the system was mostly dependent on the surface.
Measurement of the (001) surface required movement several times an hour whereas the
(012) and (110) required movement only a few times per sample (Lee et al., 2011). All analyses were conducted in the dark in order to avoid any complications associated with light exposure (Francis and Dodge, 1993; Lee et al., 2011). All of the crystals had a degree of miscut to them as evidenced by an additional reflection of the incidence beam.
The miscuts were accounted for during data reduction by summing the intensities arising from both signals to get the total intensity (Catalano et al., 2007a). Data was collected as image files using either a CCD or Pilatus detector. The CTR and RAXR data were treated and fit using MATLAB (Version 2013b, The Math Works, USA) following the methodology of Lee et al.(Lee et al., 2012) and the hematite models of Catalano et al.(Catalano et al., 2007a; Catalano et al., 2007b; Catalano, 2011; Catalano et al., 2009;
Trainor et al., 2004).
3.3 Results and Discussion
3.3.1 EXAFS
The EXAFS spectra (Figure 3.1, Table 3.1) were dominated by contributions from the backscattering of the first-shell oxygen atoms at an average distance of ~2.30 Å for both hematite sizes (Table 3.2). Pb2+ has a theoretically ideal hydration sphere of 9 oxygen atoms at a distance of 2.6 Å (Hofer and Rode, 2004). The significant difference in the lead oxygen bond distance and coordination number compared to the theoretical values
77 for the hydrated lead ion were indicative of the lead being bound directly to oxygen atoms on the hematite surface. The Pb-Fe distance of 3.3-3.4 Å was indicative of edge sharing bidentate adsorption while the Pb-Fe distance of > 3.9 Å means the lead was adsorbing in a corner sharing bidentate or a monodentate mononuclear manner (Bargar et al., 1997; Lenhart et al., 2001). The second shell Pb-Fe influence can be seen in the
Fourier Transform in Figure 1B between 2.5 and 4 Å and in the χ(k) plot where the drop in the amplitude of the third antinode was due to the Pb-Fe and Pb-O oscillations being out of phase and reducing the peak height. The LSA hematite has Pb-Fe distances of
3.32 and 3.92 Å and the HSA has Pb-Fe distances of 3.34 and 3.94 Å. The Pb-O coordination numbers were between 2 and 3 for both particle sizes which suggests that in addition to the doubly-coordinated complexes, there may be additional triply–coordinated adsorption modes. The doubly coordinated lead orientation has been suggested by several experimental papers,(Bargar et al., 1997; Lenhart et al., 2001) while the triply- coordinated lead is implicated in theoretical studies (Mason et al., 2009). Mason et al. ran a molecular dynamics simulation of lead adsorbing on the (001) surface and found that the lead would often rest on the hematite surface equidistant from all oxygen atoms
(Mason et al., 2009).
There was little difference observed in the coordination of Pb to the two different sized hematite. The coordination number for the Pb-O bond was 2.58 for the 50 nm hematite and 2.49 for the 10 nm hematite. The Pb-Fe coordination numbers were 0.64 and 0.83 for the edge and corner sharing complexes respectively on the LSA hematite and 0.82 and 0.42 on the HSA hematite. However, due to the coordination number being
78 completely correlated with the amplitude reduction factor in the EXAFS equation, there is a substantial amount of error built in to the results of the coordination number
(Newville, 2014). Overall, these results were consistent with previously published lead
EXAFS performed on hematite particles (Bargar et al., 1997; Lenhart et al., 2001).
3.3.2 (001) Surface
The (001) face, the basal surface of the hexagonal crystal structure, is considered the most stable face on natural hematite(Mackrodt et al., 1987) and thus it is the most studied hematite surface (Guo and Barnard, 2011). The (001) face ideally consists entirely of oxygen atoms doubly coordinated to iron atoms (see Figure B1)(Barron and Torrent,
1996) forming a molecularly flat surface (Catalano et al., 2009). Between a pH of 2 to
10, these oxygen atoms are singly protonated meaning the surface is neutrally charged
(Hiemstra and Van Riemsdijk, 1999) resulting in the (001) face being relatively inert
(Venema et al., 1998). The total electron density of the (001) surface in DI water shows a consistent well-defined single layer termination as evident by the distinct and rapid transition from the narrow, high electron density peaks to the much lower broader peaks of the weakly layered water (Figure 1, shown in blue). There is some disagreement in the literature as to the actual terminations of the (001) surface. Trainor et al. (Trainor et al.,
2004) studied the (001) face of a natural hematite single crystal using X-ray reflectively under humidified helium gas. They concluded that the termination of the (001) surface under these conditions was not ideal and there were in fact additional FeO6 octahedra
“adsorbed” onto the surface breaking up the ideal, flat (001) surface. Catalano (Catalano,
2011) applied in situ X-ray reflectivity to investigate the layering of water on the (001)
79 hematite surface and found that under a thin film of water, the surface was ideally terminated with a very weakly ordered water layer and significant relaxation of the Fe atoms on the top layer of hematite. The difference in crystal surface structure is likely due to sample preparation. Our sample preparation followed those of Catalano’s
(Catalano, 2011) and similar to those results, the inclusion of an additional partial layer termination into the model was not necessary to fit the data (Catalano, 2011).
There was little difference between the total surface coverage of lead at pH 4 (0.71
µmol/m2) and pH 6 (0.86 µmol/m2), which was consistent with the (001) surface remaining neutrally charged under the tested pH conditions. This surface coverage was more than an order of magnitude lower than the 13.8 µmol/m2 reported for the XPS measurements on the (001) surface by Bargar et al. (Bargar et al., 2004). There were several differences between our experiments and those conducted by Bargar et al.,(Bargar et al., 2004) notably, they used a 300-400 Å thick hematite film grown on a sapphire substrate, performed the experiments in a humidified atmosphere rather than in bulk solution, and used a higher concentration of lead in the sample solution at a pH of 7.
These differences may be important as it has also been suggested that most of the reactivity on the (001) surface is the result of defects in the crystal, either additional iron octahedra “islands” sorbed on the surface or missing octahedra in the first layer similar to the non-ideal surface detected by Trainor et al. (Trainor et al., 2004; Venema et al., 1998)
The specific makeup of the surface defects is likely dependent on the preparation of the sample prior to measurement (Lutzenkirchen et al., 2015). Performing the experiments
80 under a water film may also induce differences as it alters the relaxation of the surface atoms and the ordering of water layers (Fenter and Sturchio, 2004).
Fitting the pH 4 and pH 6 RAXR data required two lead locations. At pH 6 the near- surface peak was at a distance of 1.61 Å above the relaxed oxygen surface and at pH 4 it was at 1.5 Å. The RAXR data on the (001) surface was challenging to work with due to the low surface coverage. As a result, the edge jump of the RAXR was small and the error associated with each scan was large relative to the (012) and (110) surfaces (See
Figure C2-4 and Appendix C), negatively impacting the quality of the data and the resolution of the lead locations. Since the (001) surface in this study exhibited an ideal termination, all of the surface iron octahedra were arranged to display a face on the surface (See Figure B1). Arranging the lead atom in a bidentate structure on the (001) surface with a 2.3 Å Pb-O distance places the lead atom 1.7 Å laterally from the oxygen surface plane. This arrangement also results in Pb-Fe distances of 3.3 and 3.8 Å, close to the results of the EXAFS. The GI-EXAFS results of Bargar et al. also include an oligomeric lead complex with a Pb-Pb distance of 3.62 Å which was responsible for the high surface coverage. This complex was not detected in our EXAFS results, nor has it been reported in any other particle based Pb EXAFS studies,(Bargar et al., 1997; Lenhart et al., 2001) however, our RAXR do show additional lead locations distal to the inner- sphere adsorption in both the pH 4 and 6 samples. Analysis of the pH 6 system shows a broad second peak at ~4 Å from the hematite surface. This was consistent with the distance of a hydrated lead ion hydrogen bonded to the singly protonated hematite site,
~4.2 Å, and can therefore be assigned as such (Bargar et al., 1996). The pH 4 sample
81 shows a broad peak with a low occupancy centered at ~7.5 Å. This was too far from the surface and at too low of a surface coverage to be the oligomeric species reported by
Bargar et al. This height was also too large to be attributed to a hydrogen bonded ion. It was possible that it represents an electrostatically held ion. Lee et al, (Lee et al., 2010;
Lee et al., 2011) found a small amount of lead adsorbed to the muscovite (001) surface at pH 2 and 3.7 termed extended outer-sphere adsorption complexes at a distance of 8.34 and 9.58 Å from the surface respectively. The longer distance of the lead atom here could also be the result of the low surface coverage of Pb poorly constraining model fit locations and occupancies.
3.3.3 (012) Surface
The (012) surface, which is referred to alternatively in the literature by the equivalent
Miller indices as the (1 1 0 2), (0 1 1 2), and (1 0 1 2) or as the r-cut surface, was not as well defined as the (001) surface. Our model assumed there were only two possible terminations, both of which maintain the FeO6 octahedron structure and were oxygen terminated as the hematite was equilibrated in an aqueous solution (Catalano et al.,
2007b; Catalano et al., 2006). When each atom layer in the bulk crystal was allowed to relax independently during fitting, there was a high level of covariance between the locations of the individual atom layers in the bulk crystal (see Appendix A for more details of the model and fitting process). Therefore, we used the method of Catalano
(Catalano et al., 2007a; Catalano et al., 2007b) and allowed the bulk crystal atoms to relax as OFe, FeO and O rather than allowing each atom to individually relax as was done in the (001) surface model. Thus, the ideal termination was defined by the model as
82 FeOFeOO. Using this ideal termination, the total electron density of the (012) surface in
DI water has a high electron density peak located at z=0.9 Å (Figure 3.3). This reflects the approximate height for a subsequent iron layer and since the electron density was too high to reasonably be considered water, this peak was the likely result of an additional half layer termination at the (012) crystal surface. Thus, an additional FeO was added on to the bulk crystal model. When applying this model to the data collected under only the presence of DI Water, the half layer termination was approximately 75% of the exposed surface. The termination of the (012) surface depends upon the preparation method and the half layer may reflect the chosen annealing time and temperature (Catalano et al.,
2007a). Thus, following Catalano et al, (Catalano et al., 2007a) the surface was defined at z=0 as the midpoint between the two adjacent oxygen layers. As a result, the electron density of the bulk crystal in Figure 2 goes beyond the z=0 height.
Under the solutions tested in this experiment, the (012) surface was expected to be more reactive than the (001) surface and the RAXR results show a much higher surface coverage of Pb at pH 6, Γ = 9.63 µmol/m2, and a slightly higher surface coverage at pH 4,
Γ = 1.71 µmol/m2. The (012) surface differs from the (001) surface in that the surface oxygen atoms on the ideal termination are singly and triply coordinated to iron atoms
(Figure B.1) leading to a surface charge that exhibits greater dependence on pH. The singly coordinated and triply coordinated oxygen are arranged in a series of “ridges” and
“valleys” along the [121] direction (Barron and Torrent, 1996; Catalano et al., 2007a).
The half layer termination also provides for similar ridges and valleys, but the valleys have doubly coordinated oxygen rather than the triply coordinated oxygen of the ideal
83 termination (Catalano et al., 2007b). Previous studies on the (012) surface reported strong ordering of water bound along these ridge and valley contours of the (012) surface
(Catalano et al., 2007a). The (012) surface acts as a template for the ordering of water.
This is indicative that the surface is reactive and that the valley regions play an important role as do the ridges. Bargar et al. (Bargar et al., 2004) also investigated Pb adsorption to the (012) surface and found a slightly lower surface coverage at pH 7 (12 µmol/m2) than they reported for the (001) surface. This was still higher than our analysis showed on the (012) surface at pH 6, though not as remarkably different as the (001) surface, considering the coverage at pH 4 was only 1.71 µmol / m2.
At pH 4, the best-fit model for lead adsorption on the (012) surface resolves into 2 peaks at 1.62 and 4.08 Å above the surface with the second peak having an occupancy 25% higher than the first peak. This likely reflects the formation of similar inner-sphere complexes at the two different surface terminations. Using a Pb-O distance of 2.3 Å pursuant to our EXAFS results and those reported by others (Bargar et al., 1997; Lenhart et al., 2001), monodentate coordination at the triply coordinated valley-oxygen layer gives a surface height of 1.58 Å from an unrelaxed ideally terminated surface. The availability of the triply coordinated valley-oxygen atoms for bonding was based on these oxygen atoms bonding with water.(Catalano et al., 2007a) A similar adsorption mode is reported by Ostergren et al. (Ostergren et al., 2000) who observed a monodentate lead adsorption mode on goethite that only appeared at pH below 5. Another coordination option that fits the constraints of the EXAFS and the reflectivity involves lead bound to one triply coordinated valley-oxygen and two singly coordinated peak-oxygens. This
84 yields a lead height of 1.52 Å and corresponds to results arrived at by Mason et al.
(Mason et al., 2011) when performing a DFT analysis on lead adsorption at the (012) surface of the isostructural aluminum oxide, corundum. In their analysis, the singly coordinated oxygen atoms remained singly protonated while the triply coordinated oxygen atoms were deprotonated, allowing them to bind with the lead. The lead located at z=4.08 Å in our experiment was likely coordinated in a manner similar to the first lead location to two protonated singly coordinated oxygen and one deprotonated doubly coordinated oxygen atoms simultaneously. This arrangement yields a height of approximately 4 Å and a Pb-Fe distance of 3.22 Å, close to the EXAFS derived Pb-Fe distance of 3.3Å.
At pH 6, the surface coverage was over five times greater than at pH 4 (Table 2). There is no direct comparison in the literature, however this large jump was in agreement with the change in surface coverage observed on hematite particles (McKenzie, 1980). Lee et al. also saw a doubling of the lead coverage between pH 2 and 3.7 on the muscovite (001) surface (Lee et al., 2010). The near-surface peak shifted nearly 0.4 Å closer to the surface. This was close to the distance expected from a corner sharing bidentate complex to the triply coordinated valley-oxygens. A similar shift in lead binding to goethite from bidentate coordination to monodentate coordination with a decrease in pH was observed using EXAFS by Ostergren et al. (Ostergren et al., 2000). The largest peak was at z=3.96
Å and it accounted for over half of the total adsorption (see Table 2). This was at a similar location, but it was a five-fold increase in occupancy from the pH 4 system. This peak was likely an inner-sphere complex bound to the half layer termination. The
85 coordination was similar to the inner-sphere complex on the ideal termination layer with the lead coordinated with both the singly coordinated oxygen on the hills and the doubly coordinated oxygen in the valleys. There was also a more distant broad peak centered at z=7 Å. The distance and large rms width (1.56 Å) indicate that this was likely an outer- sphere complex (Lee et al., 2010).
3.3.4 (110) Surface
The final surface studied, the (110) surface (alternatively referred to in the literature as the (1120)), consists of equal numbers of singly, doubly and triply coordinated surface oxygen groups (Barron and Torrent, 1996). The (110) surface has only one possible termination resulting in full coordination of the FeO6 octahedra (Catalano et al., 2009).
Our model for this surface followed that used by Catalano (Catalano et al., 2009) where the surface terminates as four oxygen and one iron layer (Figure. B1 in Appendix B and table C5 in Appendix C). In this model, the outermost oxygen layer was singly coordinated, the two middle were doubly coordinated and the layer closest to the iron was triply coordinated. To avoid covariance of the oxygen layers that were only separated by
0.2 Å, the model groups the top two oxygen (singly and doubly coordinated) together and the bottom two oxygen (doubly and triply coordinated) together when allowing the bulk crystal to relax. At the same time, the iron atom was allowed to move independently
(Catalano et al., 2009). The top oxygen layer in this model for the (110) surface does not terminate at z=0, but rather 0 refers to the location of the missing iron layer and the actual top oxygen termination is at z=-0.769 Å (Catalano et al., 2009).
86 Like the (012) surface, lead coordination to the (110) surface was sensitive to changes in pH, though not to the same degree. The surface coverage was approximately 25% higher at pH 6 vs. pH 4. Of the three surfaces studied at pH 4, the (110) surface was the most receptive to lead with a surface coverage of 3.7 µmol/m2. The different surfaces likely have different pHpzc and pKa values (Venema et al., 1998) resulting in different pH- dependent surface charging and reactivity with lead.
At both pH conditions, the best-fit lead model included two sharp peaks near the surface and a broad peak more distant from the surface (Figure 4). The first sharp peak was located at z=0.6 Å or 1.37 Å above the relaxed top oxygen layer. Similar to the (012) surface, this distance indicates an inner-sphere complex, involving both triply coordinated valley-oxygen atoms and singly coordinated peak-oxygen atoms. While this was the best fit with respect to what we also know from the EXAFS, it was difficult to pinpoint the exact bonding geometry as there exist singly-, doubly-, and triply- coordinated adsorption modes that will result in a lead height of ~0.6 Å. The second peak at ~3.2 Å was nearly identical to the lattice spacing of the crystal indicating that it was possibly the same coordination on only on a different step. The broad peak at 6.2 Å was likely an outer-sphere complex as it was located at a distance consistent with that for a hydrated Pb2+ atom hydrogen bonded to the hydrated singly coordinated oxygen atoms.
The same lead locations occurred at both pH conditions, however at pH 6, the first lead peak has a slightly higher occupancy than the second.
87 3.4 Conclusion
The EXAFS and reflectivity techniques compliment each other. EXAFS gave very precise information about the local environment of the lead atom, however, since that local environment only spans a radius of a few Angstroms from the central atom
(Templeton et al., 2003) that information was somewhat limited. Based on the EXAFS, lead was bound to the hematite surface as both edge and corner sharing inner-sphere complexes in agreement with previous studies. The XR/RAXR data confirms that lead was present on all three surfaces tested at a distance indicating the presence of an inner sphere complex consistent with the EXAFS. The single crystal experiments also yielded information on the existence of an outer-sphere complex that was not visible in the
EXAFS data. When assigning coordination modes to the inner-spherically bound lead locations from the RAXR data, the EXAFS information was used to restrict the possible lead coordination states. Several possible adsorbed lead structures were eliminated from consideration due to the resulting Pb-Fe distances not conforming to the EXAFS results.
In addition to locations, the reflectivity technique also yields surface coverage, which was not directly measured with EXAFS. The XR results show that abundance of specific faces was important to the adsorption of lead on a hematite crystal. At pH 6, lead adsorbs on hematite in the order (012) > (110) > (001) and at pH 4, (110) > (012) > (001).
Adsorption on the (110) and (012) faces was more dependent on pH than on the (001) face because they have singly and triply coordinated oxygen atoms on the surface. The
(110) and (012) better match the adsorption of lead on hematite particles where the pH is a controlling variable. This was significant as the (001) surface is often considered the
88 most abundant in natural hematite (Guo and Barnard, 2011) and yet it was the least reactive face for lead adsorption at the tested pH. The results indicate that the surface topology is critical to the adsorption of lead on hematite, with the atomically flat (001) being the least reactive in agreement with the previous studies indicating that reactivity on the (001) is due to imperfections in the surface (Bargar et al., 2004; Venema et al.,
1998). The use of the two x-ray techniques together builds a more complete view of the lead adsorption process.
89
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94
10 nm Pb ( = 3.08 µmol/m2) (A) (B)
) 4
(k)
2
3 50 nm Pb ( = 3.82 µmol/m ) k (R)| (Å |
2 3 4 5 6 7 8 9 10 0 1 2 3 4 5 1 R (Å) k (Å )
Figure 3.1. Result of the EXAFS experiments. A) k3 weighted χ(k) functions and B) their Fourier transform. The data is shown with the blue circles and the fit with the red line.
95
4
pH 4 3
/Å 3
2
1 Electron Density e
0 0 2 4 6 8 10
4
pH 6 3
/Å 3
2
1 Electron Density e
0 0 2 4 6 8 10 Height above Surface (Å)
Lead Total DI water
Figure 3.2. Results of the XR / RAXR experiments for lead on the (001) surface of hematite at pH 4 and pH 6 as noted in the plots. The solid black line is the overall electron density (XR) and the red area is the lead specific electron density (RAXR). The blue dashed line is the electron density of the (001) surface in DI water.
96
4
pH 4 3
/Å 3
2
1 Electron Density e
0 0 2 4 6 8 10
4
pH 6 3
/Å 3
2
1 Electron Density e
0 0 2 4 6 8 10 Height above Surface (Å)
Lead Total DI water
Figure 3.3. Results of the XR / RAXR experiments for lead on the (012) surface of hematite at pH 4 and pH 6 as noted in the plots. The solid black line is the overall electron density (XR) and the read area is the lead specific electron density (RAXR). The blue dashed line is the electron density of the (001) surface in DI water. The tall XR peak above a height of 0 is the half layer termination.
97
4
pH 4 3
/Å 3
2
1 Electron Density e
0 0 2 4 6 8 10
4
pH 6 3
/Å 3
2
1 Electron Density e
0 0 2 4 6 8 10 Height above Surface (Å)
Lead Total DI water
Figure 3.4. Results of the XR / RAXR experiments for Lead on the (110) surface of hematite at pH 4 and pH 6 as noted in the plot. The solid black line is the overall electron density (XR) and the read area is the lead specific electron density (RAXR). The blue dashed line is the electron density of the (001) surface in DI water.
98
Table 3.1. Fits to Pb LIII EXAFS.
Particle Pb-O Pb-Fe Pb-Fe Size CN R (Å) CN R (Å) CN R (Å) 50 nm 2.58 2.31 (0.00) 0.64 3.32 (0.02) 0.83 3.92 (0.02) 10 nm 2.49 2.30 (0.01) 0.76 3.34 (0.02) 0.75 3.94 (0.02)
99
Table 3.2. Best fit model parameters from the RAXR data of lead only on the three hematite surfaces. First peak Second peak Third peak Total χ2 R-Factor Crystal pH z c µ z c µ z c µ (µmol/m2) (001) 4 4.139 0.009 1.45 (5) 0.48 (4) 0.48 (10) 7.43 (17) 0.24 (6) 0.89 (29) 0.71 (7) (001) 6 1.511 0.005 1.58 (12) 0.37 (10) 0.58 (15) 3.95 (30) 0.49 (12) 1.22 (28) 0.86 (15) (110) 4 4.835 0.061 0.60 (2) 1.38 (9) 0.22 (3) 3.09 (2) 1.53 (7) 0.25 (3) 6.28 (13) 0.8 (11) 0.82 (17) 3.71 (16) (110) 6 4.123 0.016 0.76 (1) 2.26 (4) 0.26 (2) 3.13 (3) 1.34 (5) 0.38 (3) 6.75 (9) 0.99 (10) 1.29 (16) 4.58 (13) (012) 4 2.273 0.013 1.62 (2) 0.71 (2) 0.2 (f) 4.08 (2) 1.01 (3) 0.59 (3) 1.72 (4) 100 (012) 6 3.51 0.021 1.23 (5) 0.99 (6) 0.2 (f) 3.96 (2) 5.00 (22) 0.57 (3) 6.94 (12) 3.65 (31) 1.56 (12) 9.63 (39) 2 z = height above the surface (Å); c = occupancy of the layer (µmol/m ); µ = rms width
Chapter 4: Effect of Organic Acids on Lead Adsorption on Hematite
Abstract
Two X-ray techniques, particle based extended x-ray absorption fine structure spectroscopy (EXAFS) and single crystal X-ray reflectivity (XR) were combined with batch adsorption experiments to investigate the adsorption of lead onto hematite in the presence of organic acids. Samples were run at pH 4 and 6 as the adsorption edge for lead alone on hematite rises sharply between those values for the solution conditions tested. On the hematite particles at pH 4, citric, phthalic, humic and fulvic acids enhance the adsorption of lead on hematite. At pH 6 the citrate hindered the lead adsorption while the other three acids had no impact on the amount of lead adsorbed. Both X-ray techniques showed that the lead was adsorbing partially as an inner-sphere complex, similar to how lead adsorbs to the hematite surface in the absence of organic acids. This was seen at both pH conditions, even though the presence of organic acid was necessary for lead to adsorb to the surface at pH 4. From the EXAFS data, the change in pH from 6 to 4 caused the spectra with lead and organic acids more closely resemble the lead- organic acid standards than the adsorbed lead only spectra from chapter 3. Overall, the surface structure of lead is more dependent on the surface than it is on the acid included.
The template for lead adsorption is defined by the structure of the surface and the acid serves to modify the adsorption.
101 4.1 Introduction
Lead is the most common contaminant found at superfund sites and based on its highly toxic nature it is listed second overall in hazardous substances of concern by the Agency for Toxic Substances and Disease Registry (ATSDR, 2013). Exposure to lead by infants and children can cause severe mental and physical developmental disorders (Casas and
Sordo, 2006) with a blood lead level increase from 2.4 to 30 µg/dL found to decrease IQ by 6.9 points (Lanphear et al., 2005). The inclusion of tetraethyl lead in gasoline for over
70 years and the use in paint has resulted in elevated lead levels in nearly all soils, especially in urban areas(Gooch, 2006; Mielke et al., 2011). The production of lead for these and other industrial and consumer uses produced the point source contamination at mining and production facilities resulting in its prevalence at superfund sites (USEPA,
2004).
Unfortunately, the fate and transport of this highly toxic and common pollutant has proven difficult to predict. Lead has a high affinity for metal oxide surfaces that are present in soils so it is unlikely for lead to remain in solution, however lead transport in soils has been observed in forest floors (Kaste et al., 2006). The lead in these mobile environments has been found associated with both iron oxides(Barrett et al., 2010; Kaste et al., 2006) and organic matter (Wang and Benoit, 1997). This leads to the possibility that lead is mobilized via its association with colloidal particles (Hassellov and von der
Kammer, 2008; Kalmykova et al., 2010; McCarthy and Zachara, 1989; Zachara et al.,
1994). In order to develop accurate models to efficiently and effectively remediate a contaminated waste site, it is necessary to understand the reactions of lead with the
102 surfaces that may be a vector for lead transport or sinks for lead immobilization (Bolan et al., 2014).
Hematite, α-Fe2O3, is the most thermodynamically stable of the iron oxides and as such it is commonly found in soils (Guo and Barnard, 2013). Hematite has a hexagonally close packed crystal structure and it can be represented as a series of FeO6 octahedra stacked via edge and face sharing perpendicular to the (001) plane. Hematite particles can appear with a variety of morphologies (e.g., platy, ellipsoid, spherical, etc.) which are dependent upon synthesis conditions (Penners and Koopal, 1986; Schwertmann and Cornell, 2008;
Venema et al., 1998). Macroscale hematite can likewise display a variety of morphologies, but both rarely outwardly exhibit a sharply defined crystal structure
(Cornell and Schwertmann, 2003). It is unclear what the most environmentally relevant hematite faces are with respect to adsorption processes, but it is thought that the (001),
(012) and (110) are the most common hematite surfaces (Catalano et al., 2009; Mackrodt et al., 1987).
The (001) surface is the most commonly studied hematite surface (Guo and Barnard,
2011). When crystals take on a platy morphology, the (001) face is the dominant face in terms of surface area (Colombo et al., 2012). The ideal (001) surface is atomically flat, consisting of only the faces of the hematite octahedron (see Figure B1 in Appendix B).
All of the surface oxygen atoms on the (001) surface are doubly coordinated to the underlying iron atoms. Across the range of environmentally relevant pH conditions, these oxygen atoms are singly protonated leaving the (001) surface neutrally charged
(Hiemstra and Van Riemsdijk, 1999). As such, the (001) surface is expected to be the
103 least reactive face. Catalano (Catalano, 2011) investigated the ordering of water on the
(001) face and found that the water was weakly ordered in accordance with the expected low reactivity. Bargar et al.(Bargar et al., 2004) performed grazing incidence extended
X-ray adsorption fine structure (GIEXAFS) and X-ray photoelectron spectroscopy (XPS) on the (001) surface and found that it provided an excellent sorbent surface for lead with a surface coverage of 13.8 µmol/m2 at pH 7. In Chapter 3, we used nonresonant X-ray reflectivity (XR) and resonant anomalous X-ray reflectivity (RAXR) and determined that the lead surface coverage was less than 1 µmol/m2 at both pH 4 and 6 and that the lead was adsorbed primarily in an inner sphere manner, at a distance appropriate for a bidentate corner sharing structure, and to a lesser extent as an outer-sphere complex..
Bargar attributed the high adsorption to an oligomeric Pb complex forming at the surface.
It has also been suggested that the reactivity of the (001) surface may be due to a non- ideal termination or the presence of “islands” of adsorbed iron octahedra providing surface oxygen more reactive than the doubly coordinated oxygen of the ideal termination (Lutzenkirchen et al., 2015; Trainor et al., 2004). Catalano’s (Catalano, 2011) and our experiments detailed in chapter 3 both found an ideal termination using XR.
The (012) surface, alternately referred to in the literature as the (1 1 0 2), (0 1 1 2), or (1
0 1 2) surface has a “corrugated” surface with equal parts singly and triply iron- coordinated oxygen atoms on the ideal surface termination (Figure B1 in Appendix B).
The (012) surface charge is more dependent on pH under environmentally relevant conditions due to the singly and triply coordinated oxygen atoms. Owing to the surface topology and charge, the (012) surface is expected to be more reactive than the (001)
104 surface. There are five possible terminations of the (012) surface, two of which are fully hydroxylated which is the most likely scenario when the hematite surface prepared using the methods we followed is in an aqueous system according to the results of Tanwar et al.
(Tanwar et al., 2007a; Tanwar et al., 2007b). The expression of the half layer termination is highly dependent on the preparation of the crystal surface, namely the temperature and length of time for which the crystal is annealed (Catalano et al., 2007b). The half layer termination is similar in topography to the ideal terminations, however the oxygen atoms in the valleys are doubly coordinated to iron atoms rather than triply coordinated
(Catalano et al., 2007b).
In a study of water on the (012) hematite surface, the crystal was terminated in a full layer (Catalano et al., 2007a). Due to the strong charge in the oxygen atoms, the water layer on the full termination (012) surface was well defined and followed the contours of the surface, unlike the (001) surface which had a broad near surface peak (Catalano,
2011). The half layer termination has also been found to play an important role in previous studies of arsenic onto hematite in an aqueous environment (Catalano et al.,
2007b). Bargar et al. also investigate the (012) surface using GIEXAFS and determined like the (001) surface that lead was present mostly as an oligomeric complex. In addition to the oligomeric complex, there was an inner-sphere complex directly on the surface. Combined, there was a total surface coverage of 12.0 µmol/m2 at pH 7 on the
(012) surface. In chapter 3 we found that there was a significant difference between the adsorption at pH 4 and 6 with 1.7 µmol/m2 and 9.6 µmol/m2 respectively. The lead in this case was mostly bound directly to the hematite surface and there was also an outer-
105 sphere complex. The half layer termination represented roughly 75% of the surface leading to two layers of inner sphere adsorption.
Unlike the other two surfaces studied, the (110) surface contains only one possible termination where the surface iron are all bonded with 6 oxygen atoms (Catalano et al.,
2009). The surface consists of equal numbers of singly, doubly and triply coordinated oxygen atoms arranged in four distinct layers. Like the (012) surface, the (110) surface also has a “hill and valley” topography, though not as dramatic as that on the (012) surface. The (110) surface is charged over the entire range of pH as a result of the protonation of the surface oxygen. The singly and triply coordinated surface oxygen atoms carry a partial charge and are therefore contribute more to the surface charge than the doubly coordinated oxygen which can be equalized with a single proton. The surface charge is positive at low pH and negative above the pHpzc at ~9 (Venema et al., 1998).
This surface has not been the subject of as much investigation as the (001) or (012) surfaces. Catalano reported that water was highly layered on the (110) surface and that it followed the topography of the surface (Catalano et al., 2009). As with the other two surfaces, we performed XR / RAXR on the (110) reacted with lead discussed in chapter
3. We found that the surface likely had two layers of identical termination and that lead was adsorbed as an inner sphere complex to both of those layers. There was also an additional broad outer-sphere adsorption mode present. Adsorption of lead on the (110) is dependent on the pH, with a higher surface coverage at pH 6, though the change is not as dramatic as seen on the (012) surface.
106 Structural determination studies of adsorption on hematite particles are much more common (Bargar et al., 1997; Lenhart et al., 2001; Lenhart and Honeyman, 1999) with
EXAFS being used to investigate lead adsorption on the hematite surface (Bargar et al.,
1997; Lenhart et al., 2001). The consensus of these EXAFS studies is that lead forms inner-sphere complexes with the hematite surface. The Pb-O distance is 2.3 Å and two
Pb-Fe distances are visible at 3.3 Å and 3.8 Å related to edge sharing and corner sharing surface structures, respectively. EXAFS is an element specific method as opposed to the surface specific reflectivity methods employed on single crystals, so the results are an average of the local lead environment. Lead is a particularly difficult element to perform
EXAFS on due to the high structural disorder resulting in the farthest neighbor atom distance is ~5 Å with exceptional data (Templeton et al., 2003). This constraint makes it difficult to determine the presence of outer-sphere Pb species. However, the spatial resolution of XR provides information on lead over 10 Å from the surface, which we saw in the XR results in chapter 3 (Lee et al., 2011).
Organic acids are also ubiquitous in natural systems and play a significant role in adsorption and transport processes (Tang et al., 2014). Plants and microorganisms release acids into the environment to enhance metal uptake (Hell and Stephan, 2003) or chelate harmful metals like aluminum or lead (Barone et al., 2008). As such, strong chelators like citric acid (Gao et al., 2012) or the stronger, anthropogenic EDTA (Wu et al., 2003) bind metals in solution, preventing them from adsorbing. Organic acids can also adsorb onto the surface of colloidal particles altering the surface charge of the particle surface (Vermeer et al., 1998). Acids have also been shown to lower the pH at
107 which positively charged metals will adsorb on the surface of positively charged oxides
(Wu et al., 2003). Organic acids can form ternary structures with metals. Lenhart et al.
(Lenhart et al., 2001) reported lead bridged malonate to the hematite surface. Boily et al
(Boily et al., 2005) saw the same process with phthalate bridged by cadmium to goethite forming both inner and outer-sphere structures. The acid can also act as the bridge. This is mostly reported with the high molecular weight humic and fulvic acids. Lee et al. (Lee et al., 2011) reported that fulvic acid causes lead to be distributed broadly as a function of distance from the mica surface whereas the lead was mostly adsorbed via an inner-sphere complex in the absence of fulvic acid.
In this study we performed particle (EXAFS) and surface specific (XR/ RAXR) X-ray analyses of lead in the presence of four organic acids - citric acid, phthalic acid, Elliot
Soil humic acid, and Nordic Lake fulvic acid. Citric acid was chosen because it is a common acid of known structure, found in environmental systems used by plants and microorganisms specifically to chelate lead and other metals (Barone et al., 2008; Gao et al., 2012). Phthalic acid was chosen as a recalcitrant, less soluble organic acid with a known structure to compare with citric acid. Phthalate consists of an aromatic ring with two carboxylic acid functional groups in an ortho arrangement (Hwang et al., 2007). The humic and fulvic acids were used as they represent the higher molecular weight organic acids found in the environment. We found that the adsorption of lead on hematite particles was enhanced to some degree at pH 4 by the presence of acids. At pH 6, citrate decreases the adsorption of lead and the other 3 organic acids did not have a significant influence on the amount of lead adsorbed. The adsorption of lead at pH 6 in the presence
108 of organic acids was similar in position and occupancy to the adsorption in the absence of organic acids. Broadly speaking the lead adsorption was more dependent on the surface than it was on the acid included.
4.2 Methods and materials
All solutions and suspensions were made with deionized (DI) water (Millipore 18.2 MΩ).
The citric acid was ACS grade citric acid monohydrate from Fisher Scientific. The phthalic acid was ACS grade from Alfa Aesar. The humic acid was Elliott Soil humic acid, derived from a terrestrial soil in the American Midwest. The fulvic acid used was
Nordic Lake fulvic acid, an aqueous derive fulvic acid from Norway. Both were purchased from the International Humic Substances Society.
4.2.1 Particle synthesis
Both sizes of hematite particles were prepared via forced hydrolysis methods. The low surface area (LSA) hematite was synthesized following the methods of Matijevik and
Scheiner (Matijevic and Scheiner, 1978) and Penners and Koopal (Penners and Koopal,
1986) as described in Chapter 3 and had a nominal size of 50 nm and a specific surface area of 35 m2/g. Briefly, The LSA hematite was prepared by dissolving 0.72M ferric chloride in 50 mL of 4 mM HCl and adding that to 1950 mL of HCl at 98oC. The 2 L solution was aged at 98oC for 72 hours. Upon removing from the oven, the hematite solution was immediately cooled in a water bath. NaOH was added to precipitate the hematite, which was collected and dialysed against DI water using 8-10 kDa cellulose acetate dialysis tubing (Spectrum Labs). The DI water was changed twice daily until the conductivity of the dialysate approached that of fresh DI water at which time the particles
109 were dialyised against a 10-4 M perchloric acid solution for 24 hours. The HSA hematite was synthesized following the method of Madden et al, (Madden et al., 2006), by slowly dripping 63 mL of 1 M Ferric Nitrate solution into 675 mL of boiling DI Water. When the feed solution was empty, about 45 minutes, the solution was taken off the hot plate and allowed to cool overnight on a benchtop. NaOH was added to precipitate the hematite which was removed and dialysed against DI water as described for the LSA hematite followed by dialysis against 10-4 M perchloric acid. Both particle types were verified as hematite using either Bruker or Scintag X-ray diffractometer. Particle size was determined using a Tecnai Biotwin TEM and surface area was measured via the BET method on a Flowsorb instrument (Micromeritics Instrument Corp.).
The (012) and (110) cut and polished hematite crystals were purchased from SurfaceNet
GMBH and were derived from natural sources. The (001) was acquired from Dr. Glenn
Waychunas (Lawrence Berkeley National Lab) and was also of natural origin. The crystals were prepared for the experiments following established methods (Catalano et al., 2007a) by first washing the crystals in alternating acetone and methanol baths while undergoing sonication. The crystals were then subjected to alternating NaOH and HCl baths. Following the chemical cleaning, the crystals were placed in a muffle oven at 450 oC for 4 hours to remove adventitious carbon and anneal the surface to create a smooth uniform termination.
4.2.2 Batch Adsorption
Batch adsorption experiments were performed using the hematite particles in order to evaluate the adsorption tendencies of lead on hematite necessary to inform the parameters
110 of the X-ray experiments. The batch adsorption experiments were performed by mixing lead and / or an organic acid with hematite particles in polycarbonate centrifuge tubes.
The pH was adjusted with freshly prepared CO2-free NaOH and HClO4 under a humidified nitrogen atmosphere to exclude carbonate from the system. The samples were rotated on an end-over-end rotator in the dark for 48 hours at which time the final pH was measured. The samples were centrifuged at 12000 rpm and an aliquot of the supernatant was taken to determine the surface coverage. The dissolved Pb concentration in the supernatant was measured on an inductively coupled plasma – atomic emission spectrometer (Varian Vista AX CCD-Simultaneous ICP-AES). The amount adsorbed was determined by subtracting this amount from the total known amount added to the sample.
4.2.3 EXAFS
The particle EXAFS samples were prepared by adjusting the pH of the hematite, organic acid and lead solutions independently to achieve the target pH. The three components were added in immediate succession to the background electrolyte solution and the pH was measured and adjusted if necessary. Samples were rotated on an end-over-end rotator for 24 hours in the dark to avoid any photodegradation (Borer et al., 2007; Dodge and Francis, 2002). At 24 hours, the pH was checked and adjusted if necessary and the samples returned to the end-over-end rotator for another 24 hours at which point they were removed and the final pH measured. The samples were then centrifuged down at
12000 rpm. An aliquot of the supernatant was taken for analysis to determine the surface coverage. The particle paste was placed in a PTFE (Teflon) sample holder with Kapton
111 tape windows. The samples were prepared at the Ohio State University and transported to Argonne National Lab in a cool, damp nitrogen atmosphere. The samples were kept in the refrigerator until the EXAFS measurements were performed. All manipulation of the samples occurred under a humidified nitrogen gas stream to exclude carbon dioxide.
Five standards were prepared for the EXAFS measurements. A PbO spectrum was collected by pressing a mixture of PbO and boron nitride into a pellet. The PbO was acquired from Strem Chemicals and it consisted of a blend of Pb(II) oxide structures rather than a single uniform mineral. An aqueous Pb2+ sample was collected from a
PbClO4 solution. The Pb-citrate and Pb–phthalate standards were made by mixing concentrated citric or phthalic acid with lead and adding CO2 free NaOH to the solution until a white precipitate formed. The precipitate was removed from the solution and rubbed onto a strip of Kapton tape forming a smooth uniform film. The tape was folded over on itself a sufficient number of times to yield a strong signal. The Pb-Fulvic acid standard was an aqueous mixture of lead and fulvic acid injected into a PTFE sample container with Kapton windows.
The EXAFS experiments were performed at Beamline 20-BM at the Advanced Photon
Source at Argonne National Laboratory using a Si (111) crystal monochromator and a 1 x
10 mm unfocused X-Ray beam. The beam was detuned 15% to eliminate higher order harmonics. Samples were placed in at a 45° angle to the incident beam. Sample data was collected with a 13-element Germanium detector (Canberra) in fluorescence mode that was oriented 90° to the beam. Several layers of Al foil were used to reject Fe-Kα fluorescence. Standards were collected in transmission mode using an ion chamber
112 detector purged with 100% N2 gas. The energy was varied across the Pb LIII edge at
13.035 keV. A minimum of 10 scans of each sample was taken and the scans average together to improve the signal to noise ratio. The analysis of the EXAFS data was performed with the Demeter package using ATHENA for data treatment and ARTEMIS for the model fitting(Ravel and Newville, 2005). For the model fitting, the amplitude reduction factor was set at 0.8425 for all elements in all the samples (Bargar et al., 1997;
Lenhart et al., 2001). The Debye-Waller factor was set at 0.01 for iron and oxygen shells and 0.013 or 0.008 for carbon, corresponding to a 6 member ring and a 4 member ring structure respectively(Lenhart et al., 2001). Fits were performed in R space transformed from k-space from 2.8 to as high a k value as the signal remained distinguishable from the noise using a Hanning window with a shelf of 1. The FEFF input for the model calculations was created with GaussView software. Additionally, a linear combination fitting (LCF) analysis was performed in ATHENA using the Pb–acid standards and the lead adsorbed on hematite in the absence of acid (from chapter 3) as the components.
LCF simply averages two or more known spectra together to determine the contribution of each standard to the heterogeneous sample (Scheckel and Ryan, 2004).
4.2.4 Single Crystal Reflectivity
The XR and RAXR experiments were carried out in succession, as the setup does not change. Following the annealing stage, the crystals were allowed to cool slowly in the muffle furnace until the temperature was ~70oC. The samples were then prepared by placing the crystals in a solution of lead (0.1 mM) and organic acid (1mM for citrate and phthalate and 80 mg C / L for Humic and Fulvic acids) in a 0.1 M NaClO4 background
113 electrolyte solution. The crystals were allowed to react for at least 2 hours in the dark at room temperature(Borer et al., 2007; Dodge and Francis, 2002). The crystals were then placed in a flow through holder with a Kapton window and flushed with fresh sample solution (see Figure C1 in Appendix C). The sample holder was placed on a diffractometer (Huber Ψ diffractometer at beamline 6 and Newport 6 circle at beamline
33) and aligned before scans were collected.
A complete explanation of the reflectivity methods used can be found in Appendix C.
The XR scans were collected at 12 keV while varying the incidence angle, θ and the detector arm angle, 2θ between an electron momentum transfer (q) of approximately 0.1 and 6 Å-1. This yields a lateral spatial resolution of ~0.2Å in the electron density plot.
The RAXR were collected by varying the energy across the LIII edge at a series of fixed q values. The data was collected using either a CCD detector or a Pilatus detector. Each data point was recorded as a digital image and the intensity found by integrating under the intensity of the beam reflection. Unlike mica which cleaves cleanly on the (001) surface producing an atomically smooth surface, hematite does not have a clean cleavage plane, so all of the crystals are mechanically cut and exhibited some degree of a miscut resulting in two reflections in the midzone between Bragg peaks. The miscut was accounted for by summing the intensity of both reflections when present (Catalano et al.,
2007a). The data was reduced and fit using Matlab following the methods of Lee et al.
(Lee et al., 2011).
114 4.3 Results
4.3.1 EXAFS and batch adsorption
The results of the batch adsorption experiment are shown in Figure 4.1. In agreement with previous work, lead adsorption on the LSA hematite increases sharply between pH 4 and 6 (McKenzie, 1980). Adsorption on the HSA hematite was shifted to the lower pH relative to the LSA hematite. This was in good agreement with Barton et al. (Barton et al., 2011) who saw a similar shift in the adsorption of lead with decreasing hematite particle size. The same amount of lead, 0.1mM, was added to both systems, and since the mass concentration of hematite in both systems was equal there was a larger number of surface sites available to the lead to adsorb to the the HSA hematite which may be responsible for the shift. Alternatively, the HSA hematite may have a higher percentage of sites that are preferable for lead binding. When equivalent concentrations of lead and citric acid were added to the LSA system, the adsorption edge shifted slightly to lower pH values. When the citrate concentration is 10x higher than the lead, the adsorption edge was shifted lower a full 2 pH units placing the adsorption edge between 2 and 4 on the LSA hematite. The same trend was seen with 1 mM Citrate on the HSA hematite, though the shift was not as dramatic. A similar trend is seen in related research by He et al. (in preparation) with phthalic acid, lead and hematite. The phthalic acid was added in
0.25 µM and 2.5 µM concentrations to 0.1 µM Pb and resulted in a pH shift in the adsorption edge of ~1 and 2 pH units respectively using the LSA hematite.
These results are in good agreement with previous research on organic acid - metal adsorption interactions. Christl et al. (Christl and Kretzschmar, 2001) found that copper
115 adsorption on hematite is enhanced at low pH when fulvic acid was added to the solution.
Above a pH of 6 however, the addition of fulvic acid resulted in less copper adsorption.
Lenhart et al. (Lenhart and Honeyman, 1999) found that uranium adsorption on hematite is also shifted to the lower pH values when Suwannee River humic acid is included in the system.
Based on these results, pH values of 4 and 6 were chosen for the EXAFS and reflectivity experiments under the assumption that the lead adsorbed at pH 4 interacts with the organic acid, which may or may not be the case at pH 6 as lead will adsorb to the surface in the absence of organic acids. Additionally, EXAFS samples were prepared at pH 3 on the HSA hematite due to the shift in adsorption edge as a result of particle size.
However, aside from Citrate, they were not tested due to time constraints.
4.3.2 Citrate
Figure 4.2 shows the surface coverage of lead on the hematite in the absence and presence of all of the acids from the EXAFS samples. The citrate was different from the other 3 acids in that while it enhances the lead surface coverage at low pH, it dampens the surface coverage at pH 6. Wu et al. saw this same trend when adsorbing lead onto goethite in the presence of citric acid at pH 6 (Wu et al., 2003) and it has also been observed . Citrate has been shown to mobilize lead in soils and this reduction was thus likely a result of the formation of a lead – citrate complex that keeps the lead in solution rather than adsorbing on the surface (Gao et al., 2012).
The results of the lead standard EXAFS (Figure 4.3 and Table 1) were subtly yet clearly different from the EXAFS for adsorbed lead from chapter 3 (included also in Figure 4.4).
116 The first antinode of the adsorbed lead χ(k) spectrum was located at k = 3.95 Å-1 while the first antinode of the lead-citrate standard EXAFS was located at 3.45 Å-1. The first antinodes for all of the adsorbed lead in the presence of citrate spectra (Figure 4.3) fall in between, mostly at k = 3.6 Å-1. In addition to the antinode at a lower k value, the shape of the first oscillation of the lead only was skewed to the right whereas it was more symmetric when the citric acid is added. The second oscillation in the adsorbed lead spectrum was symmetric and its amplitude was slightly higher than the amplitude of the first oscillation (when multiplied by k3). Fits to the lead only data reveal that this is because the oscillations from the two Pb-Fe distances are nearly in phase with each other and the Pb-O shell, which dominates the low k section of the EXAFS at the second anti- node (shown in Figure 4.6). The two Pb-O distances present in the adsorbed lead citrate spectrum, however, were out of phase at the second antinode causing the amplitude to be reduced. This was also reflected in the Fourier transform and in the corresponding model fit. The adsorbed lead samples required only one Pb-O distance and two Pb-Fe distances to fit the data well (Chapter 3). When fitting EXAFS for adsorbed lead in the presence of citric acid an additional Pb-O bond was required. This was clear in the Fourier transform as the Pb-O shell was shorter and wider than in the lead only sample. There was also an additional shell corresponding to a Pb-C distance when citrate was present.
Unfortunately, the signal from carbon was weak as it is a low Z element and it is difficult to distinguish carbon from oxygen (Nelson and Miller, 2012).
In applying the LCA analysis to the EXAFS for lead adsorbed on hematite in the presence of citrate, the corresponding EXAFS spectra for the adsorption of lead on HSA
117 or LSA hematite, Pb–citrate, and Pb2+ were used (Figure 4.4, Table 4.2). At pH 6, the
EXAFS for the adsorption of lead in the presence of citrate onto LSA hematite was determined to be comprised of roughly equal part lead citrate and adsorbed lead with
5.2% Pb2+. At pH 4, the balance shifts to 64.8% lead citrate and 21.1% adsorbed lead with the remaining 14.1% made up by aqueous lead. The HSA hematite samples show a similar pattern with the adsorbed Pb decreasing from 54.9% to 33.6% and 26.9% at pH 6,
4 and 3 respectively. In response to this shift, the lead-citrate contributions increased from 32.3% to 56.1% at pH 6 and 4 and remaining at 56.1% at pH 3. The aqueous lead made up the balance. As expected, this analysis shows that the influence of the organic acid increases as the pH decreases below the lead only adsorption edge. It also shows that the lead was likely still adsorbing directly to the hematite surface in a reduced amount even at pH 4.
Modeling the EXAFS data yields a best-fit Pb-O distance for a lead with citrate sample of
~2.3 Å, which was in similar to the results of our lead-only system (Chapter 3) and previously published lead only adsorption data (Bargar et al., 1997; Lenhart et al., 2001).
In the presence of citrate this best-fit Pb-O distance remains fairly constant at both pH 4 and 6. The second Pb-O distance, which was not present in the adsorbed lead only
EXAFS, was longer, between 2.47 and 2.50 Å for all samples. Kourgiantakis et al
(Kourgiantakis et al., 2000) synthesized a lead-citrate crystal by slowly precipitating a 2:1 mixture of citrate and lead, and analyzed it via X-ray diffraction. They determined there are four Pb-O bonds ranging in distance from 2.39 to 2.53 Å. As the Pb-O distances in the Pb-Citrate complex are longer than our adsorbed lead Pb-O distance and shorter than
118 the Pb-O distance of hydrated Pb2+(Hofer and Rode, 2004), we therefore attribute the longer distance to direct interactions with citrate. As the pH decreased in the systems with HSA hematite from 6 to 3, there was an increase in the Pb-C distance and both Pb-O distances (Table 4.1). This was interpreted as resulting from a shift from lead being bound directly to the hematite surface to an increase in the interaction dependent on the presence of citrate or possibly an outer-sphere complex Pb-Citrate complex. Aqueous
Pb2+ has an ideal primary hydration shell consisting of 9 oxygen at ~2.6 Å (Hofer and
Rode, 2004). While lead and hematite were both positively charged at pH 3, the citrate adsorbed on the surface acts to make the surface charge more negative and therefore more electrostatically favorable for the adsorption of metal cations like lead (Simanova et al., 2011).
The results of the lead – citrate XR and RAXR experiments are shown in Figure 4.6 along with the results from the lead only systems examined in chapter 3. The presence of citrate increased the amount of lead adsorbed relative to the lead only system in all cases except the (012) surface at pH 6. On the (012) surface, which adsorbed more lead than the (001) or (110) surfaces (chapter 3), the lead surface coverage was reduced to less than half what it was with lead alone (see Figure 4.6 and Table 4.3). The total lead surface coverage on the (110) and (012) surfaces remained relatively constant with the pH change, with changes of less than 5% (Table 4.3). The lead locations also remained fairly constant, however, there was a shift in the relative occupancy between the lead locations.
At pH 4, the lead locations shift away from the surface relative to pH 6. This was likely because at pH 6, the lead was adsorbing to the hematite surface as it would in the absence
119 of citrate as was seen with lead in a biofilm on the hematite surface (Templeton et al.,
2001). At pH 4, the lead adsorption was more dependent on the presence of citrate. The broadness of the third peak indicates that it was an outer-sphere complex. It was possible that it was either a lead ion being held electrostatically as the adsorbed citrate will lower the surface charge and the electrostatic repulsion. The other option involves the association formation of a lead citrate species as an outer-sphere complex, which has been seen with lead and EDTA (Bargar et al., 1999). From Chapter 2, we know that citrate adsorbs onto the hematite particle as both an outer- and inner-sphere complexes making either of these bonding modes possible. This was not the case for the (001) surface, which was the least reactive of the 3 surfaces. There was an approximately 3- fold increase in the amount of lead adsorbed when the citrate was added to the system at both pH values tested (Table 4.3). At pH 6, there was ~30% more lead adsorbed than at pH 4 and the lead had a broader distribution and higher average height. This was however the most difficult surface to measure and the errors were much higher.
4.3.3 Phthalate
Phthalate is a weak dicarboxylic acid comprised of an aromatic ring with two acid groups in an ortho arrangement. It has a lower solubility than straight chain tri-basic citrate, but it is still reactive with lead according to NIST 46.6 in visual Minteq 3.0 (see Pb –
phthalate speciation in appendix B, Figure B.3) (Gustafsson, 2011). Phthalate is known to chelate metals and is capable of forming long chain polymers or crystals with lead and other metals(Baca et al., 2004; Marandi et al., 2007). Boily et al. (Boily et al., 2005) found that phthalate formed ternary complexes with cadmium on the hematite surface
120 where the cadmium formed a bridge between the hematite and the phthalate. The phthalate adsorbed onto the surface cadmium as both inner-sphere and outer-sphere species. Hwang et al. found that phthalate by itself on the hematite surface forms predominantly inner-sphere complexes at low pH and outer sphere complexes under circumneutral pH conditions (Hwang et al., 2007).
Like the citrate EXAFS, the presence of phthalate causes the first antinode in χ(k) to shift to lower k values (Figure 4.8). However, the effect was not equal at pH 6 and pH 4. The
EXAFS for both size hematite particles show the shift in the first antinode to lower k values for the pH 6 samples less than that for the pH 4 samples. For the pH 4 samples, the amplitude of the second oscillation also decreases in a manner that was similar to that observed in the EXAFS for the lead citrate standard. The amplitude of the second oscillation in the EXAFS for the pH 6 samples stay relatively constant, which was similar to that observed for the EXAFS of lead only adsorption on hematite (see Figure
4.3). This was interpreted as the lead adsorbing directly to the hematite surface and not significantly under the influence of phthalate. He et al. (in preparation) showed that like citrate, phthalate will shift the lead adsorption edge to lower pH values. Thus at pH 4, the adsorption of lead was more dependent on the presence of phthalate to adsorb onto the hematite surface than it was at pH 6. The existence of a lead – phthalate polymer
(Marandi et al., 2007) was not evident in the EXAFS as the Pb-C bond length was too long. The polymer form has a four member Pb–carboxylate group which results in a Pb-
C distance ~2.9Å whereas the model fits (Table 4.1) yield a Pb-C distance over 3.2 Å for
121 all samples. This was consistent with the existence of six member or greater ring structure (Lenhart et al., 2001).
The LCA analysis (Figure 4.6, Table 4.2) indicates the adsorbed lead contributions to the
EXAFS for lead adsorption in the presence of phthalate was 78.5% for LSA and 74.5% for HSA at pH 6. Contributions from the lead–phthalate standard and aqueous Pb2+ make up the balance with the fits requiring slightly more aqueous Pb2+ than lead-phthalate for both hematite particles. When the pH was reduced to 4, the contribution from the adsorbed lead EXAFS decreased to 54.1 and 50.1% on the LSA and HSA hematite particles, respectively. The lead phthalate contribution correspondingly increased to 26.9 and 36.1% on the LSA and HSA with the remainder comprised of aqueous Pb2+ contributions. While the fraction attributed to the lead–phthalate standard increased as the pH dropped to pH 4, over half of the spectrum was still attributed to the adsorbed lead on both hematite particles sizes.
The lead–phthalate XR and RAXR results are shown in Figure 4.8 and Table 4.3. On the
(001) surface, the addition of phthalate had little impact, with respect to the lead only system, on the amount or location of the lead at either pH. The lead surface coverage was only slightly increased in the presence of phthalate, though not outside the margin of error. The surface coverage of lead was lower on the (001) surface than on the (110) and
(012) surfaces and it was present in a broad distribution that slightly favored inner sphere complexation. Contributions associated with an outer-sphere complex were also evident further from the surface, which was similar to that for results collected for lead adsorption on the (001) surface in the absence of acids from chapter 3.
122 The surface charge arising from the protonation state of the (110) and (012) surfaces are more pH dependent than that at the (001) surface (Venema et al., 1998), which was consistent with the observation in Chapter 3 of higher lead surface coverage at pH 6 as compared to pH 4. The presence of phthalate caused a decrease in the lead surface coverage on the (012) surface at pH 6 (see Table 4.4) when compared to the lead only system. The distribution of lead at the interface did not significantly deviate from the lead only electron density profile showing two inner-sphere complexes on the two terminations (peaks at 1.62 and 4.00 Å) and a broad distribution further from the surface
(Figure 4.9). The main contribution to the reduction in surface coverage at pH 6 likely reflects phthalate bonding with the lead in solution, forming a complex preventing it from binding to the (012) surface. At pH 4, there was a slight decrease in the lead surface coverage compared to the lead only at pH 4, though the difference was still within the margin of error. Similar to the lead only at pH 4, only two lead locations were necessary to fit the data. The results show the first lead location was closer to the surface, at 0.92 Å in the presence of phthalic acid, however quality scans were not taken to a high enough q to provide the same level of precision in the lead distribution. Overall, lead coordination to the (012) surface appeared to involve contributions from an inner-sphere Pb complex and electrostatically-bound Pb.
The adsorption of lead in the presence of phthalic acid on the (110) surface at pH 6 was the only acid-containing system that showed a significant increase in lead adsorption, as it increased from 4.58 µmol / m2 in the absence of phthalic acid to 6.91 µmol / m2 in the presence of phthalic acid. This increase was mostly associated with the second peak and
123 the broad distribution while the occupancy of the lowest z peak was decreased. At pH 4, the outer-sphere complex on the (110) surface exhibited lower occupancy relative to both the lead-only system at pH 4 and the lead–Phthalate system at pH 6. This could be the result of phthalate adsorbing onto the inner sphere lead forming a ternary lead bridging complex as described by Boily et al (Boily et al., 2005) with cadmium. This would sterically hinder lead from adsorbing in an outer-sphere manner.
4.3.4 Fulvic Acid
Like the EXAFS for the Pb-citric acid standard, the EXAFS for the Pb-fulvic acid standard deviated significantly from EXAFS for lead adsorbed to hematite in the presence (Figure 4.10) and absence(Figure 4.3) of fulvic acid. The first oscillation in the
Pb–fulvic acid standard EXAFS was symmetric and present at a lower k than the corresponding oscillation in the lead-only EXAFS (Figure 4.3). Each succeeding antinode decreased in intensity with increasing k, which differed to that for the lead-only
EXAFS that had an asymmetric first oscillation and peak c for the second antinode that was higher than the first. When fulvic acid was included in the hematite lead systems at pH 6, the EXAFS spectra resembled that for the lead only EXAFS (Figure 4.3) more than the EXAFS for the lead – Fulvic standard (Figure 4.3). The model fit was also closer to the lead only model with only one oxygen shell necessary to fit the data. The best-fit Pb-
C distances of 3.08 Å and 3.04 Å on the HSA and LSA hematite, respectively, was shorter for systems with fulvic acid than for the other acids which suggests a shift to four membered ring structures as apposed to 6 or larger member ring structure as was seen with the other acids (Bargar et al., 1997; Lenhart et al., 2001). At pH 4, the EXAFS look
124 more similar to the Pb-Fulvic than the Pb only. On both sizes of hematite two oxygen distances are again necessary to fit the data.
The LCA fitting (figure 4.5) for the systems with fulvic acid showed similarities to those for systems with phthalic acid, with contributions associated with EXAFS from adsorbed lead being more significant than those for Pb-fulvic acid. At pH 6, over 80% of the
EXAFS for lead adsorbed to hematite in the presence of fulvic acid for both hematite particle sizes was attributed to EXAFS for adsorbed lead with the Pb–fulvic standard making up most of the rest of the balance. On the HSA particles at pH 4, 63.3% of the
EXAFS was attributed to adsorbed lead with 21.7% and 15% related to Pb-fulvic and aqueous Pb2+, respectively. The LSA hematite was much more similar to the Pb-fulvic spectrum than was the HSA hematite at 39.4% vs. 21.7%. Contributions from adsorbed lead EXAFS made up the majority of the remainder at 56.1%, with Pb2+ contributions at just 4.5%.
An alternate model was used to fit the RAXR of fulvic acid (Figure 4.12). In it, the
Gaussian peak shape was replaced in the last peak with a series of overlapping Gaussian peaks that get smaller each successive peak resulting a distribution skewed away from the surface (Lee et al., 2008). This approach has been shown to more accurately represent the electron distribution of fulvic acid adsorbed on muscovite (Lee et al., 2008; Lee et al.,
2011; Lee et al., 2007). The overall lead surface coverage was higher when fulvic acid was present for all systems except that for the (012) surface at pH 6. On the (001) surface, the surface coverage was increased four fold versus no acid system and there was a slight dependence on pH. At both pH conditions, the lead was adsorbed in broad
125 distribution, more so at pH 4, than at pH 6 where distinct peaks were apparent. This was similar to the effect observed for the other acids on this surface (see Figures 4.7 and 4.9).
The (110) surface showed a proportional increase in the lead coverage in the presence of fulvic acid. At pH 4, the lead locations were very close to those in the lead only system and followed the same pattern of two near-surface sharp peaks and a more distant broad peak. At pH 6, the lead peak closest to the (110) surface had only one quarter the occupancy of that for the same surface evaluated in the absence of organic acids. The occupancy of the second peak was about the same, however, a big increase was observed in the third diffuse layer from 1 µmol/m2 to 5.87µmol / m2. The latter represented 75% of the total lead adsorption in the presence of fulvic acid and 28% more than the total lead adsorbs in the absence of organic acids at pH 6. Lead adsorption on the (012) surface was hindered by the presence of fulvic acid at pH 6 and enhanced at pH 4. In both cases, the average lead location was further from the surface than in the lead only system. At pH
4 and 6, outer-sphere lead represented 75 and 60%, respectively, of the total adsorption.
Fu et al. (Fu and Quan, 2006) observed fulvic acid forming an outer-sphere complex with hematite at pH 4 and undergoing both specific and electrostatic adsorption at higher pH.
Between pH 5 and 6 the hematite surface became negatively charged as a result of the fulvic acid (Fu and Quan, 2006). Lee et al.(Lee et al., 2011) investigated the impact of fulvic acid on lead adsorption to the muscovite (001) surface and found a similar pattern where the lead simultaneously adsorbed to the surface in an inner sphere manner and appeared to be embedded in a fulvic acid film.
126 4.3.5 Humic acid
The EXAFS for lead adsorbed to hematite in the presence of humic acid (Figure 4.13) follow the same pattern as the other samples. The first antinode was at a higher k values at pH 6 than at pH 4 and the second antinode is higher at pH 6 than at pH 4. This implies at pH 6 that the lead was bound in a manner very similar to that for lead adsorption in the absence of organic acids at pH 6. Fitting the data for the LSA hematite revealed that the signal from carbon was not apparent, however the longer oxygen distance was present.
The adsorption of lead in the presence of humic acid on the HSA hematite at pH 6 showed a lead – carbon interaction and the distance for this was 3.09 Å indicating a shift to a four member ring. This is consistent with the work of Orsetti et al. (Orsetti et al.,
2006) where lead was found to form a bridging complex between humic acid and the goethite surface. At pH 4, there was again a shift in the χ(k) towards a spectrum that looks more similar to the Pb-acid standards. The LCF analysis (Figure 4.5) attributed
87.1% of the HSA pH 6 spectrum to the adsorbed lead spectrum versus 63.3% at pH 4.
In comparison, the corresponding spectrum for LSA hematite showed contributions from the adsorbed lead EXAFS of 86.2 and 56.5% at pH 6 and 4, respectively. Most of the balance was made up by the Pb-fulvic standard (Figure 4.5). Similar to the fulvic acid systems, EXAFS from hydrated Pb2+ represented only a small fraction of the spectrum at pH 6 indicating that little adsorbed lead was held electrostatically as an outer-sphere complex. At pH 4, the hydrated Pb2+ share was increased to a little more than 10.2 and
11.5% on the LSA and HSA hematite respectively. No Pb-humic acid standard was collected, so the LCF was performed with both fulvic acid and phthalic acid as well as
Pb2+. Fulvic and phthalic acids were chosen to use in place of humic acid as humic and 127 fulvic acid are both large heterogeneous molecules and humic and phthalic both has aromatic rings.
The humic acid chosen, Elliot Soil humic acid is derived from terrestrial soil and it is not very soluble at low pH (Ritchie and Perdue, 2003; Thurman and Malcolm, 1981). No reflectivity data was collected at pH 4 as the humic acid quickly precipitated out of solution at that pH and ionic strength. Thus, the influence of humic acid on lead adsorption to the single crystals was only tested at pH 6 (Figure 4.13). At pH 6 in a 0.1
M NaClO4 solution, the humic acid stayed in solution for several days, but it slowly precipitated when mixed with the 0.1 mM Pb solution. This occurred in the 0.1 M
NaClO4 background solution as the pH was held constant indicating that the lead and humic acid do directly interact. The RAXR shows on the (110) and (012) surfaces that the humic acid inhibits the adsorption of lead at pH 6 (Figure 4.13) as compared to lead only. The positions of the lead atoms on these surfaces in the presence of humic acid did not change significantly from the lead only case, however the distribution did. On the
(012) surface, all three peaks had similar occupancies resulting in a more even distribution as opposed to the lead only case where the majority was present as an inner sphere complex on the half layer termination. On the (110) surface, the distribution was closer to the surface than the lead only, though the closest inner-sphere complex had a lower occupancy. The (001) face featured a low occupancy broad distribution of lead.
The reduction in surface coverage on the (012) and (110) surfaces was likely partially due to the humic acid chelating the lead, preventing it from reaching the surface. This was reported by Yan and Bai (Yan and Bai, 2005) when they adsorbed humic acid and lead to
128 chitosan hydrogel beads at pH 6.5. This is not a similar surface to hematite, however, when the lead and humic acid were added at the same time, the lead surface coverage decreased compared to the scenario when the beads were allowed to equilibrate with one additive at a time. Alternatively, the humic acid may be precipitating onto the hematite surface blocking the lead from reaching the surface. AFM investigations of NOM on the hematite surface has shown that NOM can form micelle like aggregates on the surface in the presence of groundwater at pH 4 (Namjesnik-Dejanovic and Maurice, 2001). At pH
6, humic acid is more soluble than at pH 4, however the Elliot Soil humic acid used in this experiment did precipitate at pH 6 in the presence of lead.
4.3.6 General Discussion
EXAFS modeling has difficulty with heterogeneous systems (Templeton et al., 2003).
While it provides very precise information, the method determines an average all of the atoms in the local environment. This works against the method when analyzing samples with multiple similar local environments as the difference between the different bonding environments will be lost (Nelson and Miller, 2012). Using lead as a target atom is especially difficult due to its high structural disorder resulting in a poor signal to noise ratio at higher k and as a result even in homogenous samples it is difficult to get good data above 10 k when targeting lead (Templeton et al., 2003). For this study, the best-fit models to the lead EXAFS with organic acids in the system all yielded similar results.
The lead was adsorbed directly to the surface with edge-sharing and corner-sharing Pb-Fe distances for all systems studied. In the lead only system from Chapter 3, the data is fit well with only one oxygen shell with a distance of 2.3 Å and a coordination number
129 around 2.5, which is in agreement with previous studies (Bargar et al., 1997; Lenhart et al., 2001). When an acid was added, an additional Pb-O shell was necessary to fit the data. The additional Pb-O shell was slightly out of phase with the original Pb-O resulting in an increase in the amplitude of the first oscillation in the EXAFS and a decrease in the amplitude of the second oscillation (see Figure 4.6). While the first two Pb-O shells were very close together, the inclusion of both Pb-O distances was justified because the system was complex and multiple lead – acid coordinations were possible as was previously demonstrated with lead – acid (Kourgiantakis et al., 2000) and lead-phthalate complexes
(Marandi et al., 2007).
At pH 4, the EXAFS provide more evidence of interactions involving the organic acids whereas at pH 6 the EXAFS were more similar to those for the lead only EXAFS. This suggests that lead forms a bridging complex between the hematite surface and the organic acid. There is evidence in the literature of metals forming bridging complexes with ligands when the metal adsorption is enhanced by the presence of the ligand. Boily et al.
(Boily et al., 2005) used EXAFS in their study of phthalate and cadmium(II) on the iron hydroxide goethite and found a similar scenario where the Cd was bound to the goethite surface with phthalate adsorbed on top of the Cd. Lenhart et al. (Lenhart et al., 2001) also found a lead bridging complex between malonate and the hematite surface. In it, the malonate formed both four and six membered rings with the lead as evidenced by the Pb-
C distances being between the two distances expected for four and six membered rings.
Redden et al. found citrate involved in a uranyl bridging complex on the goethite surface
(Redden et al., 2001). Based on the Pb-C distance (Table 4.1), the citrate and phthalate
130 both formed bidentate complexes with the lead resulting in 6 or greater membered rings.
For fulvic and humic acids at pH 6 the best-fit to the EXAFS suggests complexes were formed that comprised both two carboxylic groups or one carboxylic group as the Pb-C distance was between the distances expected for those structures. According to Lenhart et al. (Lenhart et al., 2001) a six membered ring yields a Pb-C distance of about 3.2 Å whereas s four membered ring yields a corresponding Pb-C distance about 2.9Å.
Ostergren et al. (Ostergren et al., 2000) also found evidence of a ternary complex involving lead forming a bridge for sulfate on the goethite surface that resulted in enhanced lead adsorption. They credited the formation of hydrogen bonds between the
SO4 and the surface with having a stabilizing effect on the corner sharing Pb surface complex.
Simanova et al. investigated metal – oxalate complex adsorption at the goethite particle – water interface with EXAFS and FTIR and found many of the metals tested formed ternary surface complexes (Simanova et al., 2011). Of the several metal-oxalate complexes studied, the one with cobalt was adsorbed on the surface as an intact outer- sphere complex before becoming a ternary complex. The charge of the Co(II) ion was reduced by the combination with the anion, making electrostatics the driving force of the adsorption process rather than hindering it as it would for the Co(II) ion (Simanova et al.,
2011). This pathway is one potential mechanism for the lead adsorbing directly to the surface under conditions (i.e., low pH) where the lead is not normally adsorbed. In this scenario the order of addition of the components may matter. For the EXAFS and batch adsorption experiments, the hematite was equilibrated at the target pH in a salt solution
131 and the acid and lead added to the hematite in immediate succession. The RAXR differed in that the acid and lead stock solution were both at the target pH and ionic strength when they were diluted in the NaClO4 solution. The hematite crystal was added last and the available surface area was around 5 orders of magnitude less than the particle experiments. Order of addition should be less important for citrate, which adsorbs primarily, though not entirely as an outer-sphere complex (chapter 2), than for phthalate which adsorbs primarily as an inner-sphere complex at low pH (Hwang et al., 2007).
Looking at humic acids, Reiller et al. found that when the humic acid was added first to a goethite solution and allowed to equilibrate, the cadmium surface coverage was less than when they were added simultaneously or cadmium was added first (Reiller et al., 2005).
Yan and Bai (Yan and Bai, 2005) also saw a dependence on order of addition of lead on chitosan hydrogel particles where less lead was adsorbed if both humic acid and lead were added together rather that one at a time, independent of order. Davis and Bhatnagar
(Davis and Bhatnagar, 1995) observed that when cadmium is added to a hematite solution before humic acid, the adsorption is less than if the order is reversed or the Cd and humic acid were added simultaneously. They attributed the difference to the humic acid adsorbing to the hematite surface, blocking the cadmium from forming inner-sphere complexes. Order of addition was beyond the scope of the present work, however it is a line of questioning that bears more investigation.
The data collected from the X-ray reflectivity experiments provide additional insight into lead coordination to hematite in the presence of organic acids. Focusing specifically on the surface rather than the target element yields a larger field of view and gives new
132 information about the adsorption mechanisms of lead on hematite surfaces. Overall, based on the observed similarities in the electron density profiles of the different acid additions and the lead only profiles on the same surface, adsorption appeared to depend more on the exposed crystalline surface than on the acid included in the system. This was particularly the case at pH 6. In essence, the exposed hematite surface appears to act as a template for lead adsorption and the acids modify that template. This was similar to the results collected by Catalano et al. (Catalano et al., 2007a; Catalano, 2011; Catalano et al., 2009) of water layering on (001), (110) and (012) hematite surfaces. They found that the water layering depended on the structure of the surface with the water adsorbing more strongly on the corrugated (012) and (110) surfaces than on the (001) surface in a manner that followed the respective surface contours. On the (001) surface, the water adsorbed weakly to the surface in one layer, mimicking the flat surface. This was also a point of agreement between the EXAFS results and those collected with RAXR. From the EXAFS results at pH 6 it appears that lead coordination to hematite for the lead-acid systems was more similar to the lead-only system than the lead – acid standards. The
LCF analysis indicates for all samples with organic acids that at least half of the lead
EXAFS spectrum was attributed to the adsorbed lead at pH 6 (See Figure 4.5). At pH 4, however, the lead was mainly adsorbed due to the simultaneous presence of adsorbed acid (Figure 4.1). Only citrate showed a acid – lead complex being responsible for more than 50% of the EXAFS. All of the other systems show more similarity to the lead only spectrum than the lead–acid spectrum
On the (001) surface, the RAXR results showed lead located in a broad distribution
133 relative to the hematite surface. The citric and fulvic acids result in higher lead adsorption than the humic and phthalic acids at both pH 4 and 6. With citric and fulvic acids, the RAXR results suggest the lead was absorbed to the (001) surface in the form lead–acid film. Lee et al. found a similar scenario of lead with fulvic acid on the (001) muscovite surface at pH 3.7 (Lee et al., 2011). Templeton et al. looked at the impact of a biofilm on the (001) surface of hematite and the (001) and (012) surfaces of corundum using X-ray standing wave and determined that the biofilm did little to alter the lead adsorption mechanisms. The biofilm was deposited on the crystal surface by
Burkholderia cepacia which were metabolically inactivated by exposure to X-rays before being equilibrated with lead. In all cases, the lead still preferentially adsorbed directly to the surface (Templeton et al., 2001).
At pH 6, all of the systems on the (012) surface except for the one with citrate followed the same distribution pattern set by the lead only system. Two inner sphere complexes, one on each of the two surface terminations and a broadly distributed outer sphere complex (Figure 4.7). With citrate at pH 6, the outer-sphere peak was not as broad and had a lower occupancy. At pH 4, the influence of citrate was again different from that attributed to the fulvic and phthalic acids in that there was a large lead occupancy in the outer-sphere peak whereas the systems with phthalic and fulvic acid both had broader, more even distributions. The citrate adsorption can be attributed to the adsorption of an intact lead-citrate complex via electrostatic interactions. This was in agreement with the
EXAFS LCF analysis, which showed that the influence of citrate was different from that due to the other acids (Figure 4.3). An intact outer-sphere complex was also observed in
134 the adsorption of lead onto goethite surface in the presence of EDTA (Bargar et al.,
1999).
Lead adsorption on the (110) surface was divided into 2 sharp peaks near the surface for all of the acids tested (Figures 4.7, 4.9, 4.11, and 4.13). This was attributed to inner- sphere peaks on two different, but identical termination layers, similar to the lead-only scenario. For all systems, save for that with citrate at pH 6, there was also a broad adsorption peak further from the surface. Similar to the influence of citrate on lead adsorption to the (012) surface, the third peak was smaller and more well defined than was the case for either the lead only system or the systems with the other acids. Both pH
6 citrate systems, on the (110) and (012), may be attributable to the formation of a hematite – Pb – citrate bridging complex.
The (001) surface is considered the most common hematite face, however it may not be the most important for adsorption reactions. The (012) surface, commonly considered the second most abundant face shows a much higher reactivity and matches the particle adsorption curves much better than the (001) with the sharp rise in lead adsorption between pH 4 and 6. The general trend of the surfaces in this study was that the more topographically variable the face was, the more reactive it was. This is important for particle studies not only because it shows that the morphology may be very important, but also because it can be extrapolated that higher rate of defect sites results in increased reactivity by introducing a higher density of singly and triply coordinated surface oxygen atoms, which matches previous research (Venema et al., 1998).
135 Despite marked differences in amount of lead adsorbed to the different surfaces, with differences of up to nearly an order of magnitude, there was relatively little difference in lead coordination for the two particle sizes tested. The LCF analysis shows that the
EXAFS spectra of lead adsorbed on the HSA particles in the presence of acids was slightly more similar to the lead adsorbed than to the lead–acid than it was for the LSA particles. An interrogation of the EXAFS spectra agree with this as evidenced by the location of the first anti-node, which was moved to lower k values in the presence of an acid (see Figures 4.3, 4.4, 4.8, 4.10, 4.12). The difference was also seen in the relative amplitude of the first two oscillations where the amplitude of the second anti-node was lower in the presence of organic acids (Figure 4.3). The particles express the same amorphous morphology (Figure A1 in Appendix A), which may explain the similarity in lead coordination despite the size difference. There was a shift in the adsorption edge to lower pH with the HSA hematite, but coordination did not change much with particle size
(Table 3.1 and 4.1). Thus, it seems the shift in adsorption was related to additional sites of the same type being available. Madden et al. used 7, 25 and 88 nm diameter hematite particles to investigate the size dependency of copper(II) adsorption on hematite and found that the edge was shifted to lower pH values partially as a result of an increase the number of sites associated with particle edges or other topographic features (Madden et al., 2006). More research on more variable sizes and morphologies of hematite is necessary to draw any firmer conclusions about the effect of particle size.
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144
100
90
80
70
60 Lead only LSA 0.1 mM Citrate LSA 1 mM Citrate LSA 50 Lead only HSA 1 mM Citrate HSA 40 Percent adsorbed
30
20
10
0 1 2 3 4 5 6 7 8 9 10 11 pH
Figure 4.1. Surface coverage of lead on LSA and HSA hematite particles in the presence and absence of citric acid.
145 100
90
80
70
60
50
% adsorbed 40
30
20 Lead only Citrate 10 Phthalate Fulvic acid Humic acid 0 3 3.5 4 4.5 5 5.5 6 pH Figure 4.2. Lead adsorption results for the EXAFS samples on the HSA hematite
146 Pb2+
Pb Phthalate
Pb Fulvate ) 4 − (k)
Pb Citrate 3 k (R)| (Å |
LSA Pb (0.7 mM) Only pH 6
HSA Pb (1 mM) Only pH 6
2 3 4 5 6 7 8 9 10 0 1 2 3 4 5 −1 R (Å) k (Å )
Figure 4.3. EXAFS of standards used for the LCF. The vertical line at k = 3.95 Å-1 is the antinode location of the lead only adsorbed on hematite and the vertical line at k = 3.45 Å-1 is the location of the Pb-Citrate antinode. Data in blue, fit in red.
147 LSA Pb (0.7 mM) − Citric acid (7 mM) pH 6
LSA Pb (0.7 mM) − Citric acid (7 mM) pH 4 ) 4
HSA Pb (1 mM ) − Citric acid (10 mM) pH 6 − (k) 3 k HSA Pb (1 mM) − Citric acid (10 mM) pH 4 (R)| (Å |
HSA Pb (1 mM) − Citric acid (10 mM) pH 3
2 3 4 5 6 7 8 9 10 0 1 2 3 4 5 −1 R (Å) k (Å )
Figure 4.4. EXAFS of citrate and lead adsorbed on LSA and HSA hematite as a function of pH. Data in blue, fit in red. The vertical line at k = 3.95 Å-1 is the antinode location of the lead only adsorbed on hematite and the vertical line at k = 3.45 Å-1 is the location of the Pb-Citrate antinode. The different lead concentrations were used to maximize surface coverage of lead.
148
Adsorbed Pb Pb−Citrate Pb−Fulvic LSA HSA Pb−Phthalate Pb2+ 100
80
60 Percent
40
20
0 pH4 pH6 pH4 pH6 pH4 pH6 pH4 pH6 pH3 pH4 pH6 pH4 pH6 pH4 pH6 pH4 pH6 Citrate Phthalate Fulvic Humic Citrate Phthalate Fulvic Humic
Figure 4.5. Results of the Linear Combination fit to the Pb EXAFS.
149 1.5 exp (A) O 1 Fe1 Fe2 total 0.5
(k) 0 3 k −0.5
−1
−1.5 2 3 4 5 6 7 8 9 10
1.5 exp O1 (B) O2 1 C Fe1 Fe2 0.5 total (k)
3 0 k
−0.5
−1 2 3 4 5 6 7 8 9 10 1 k (Å )
Figure 4.6. EXAFS data, fit and contribution from each path from (A) lead only and (B) Lead and Fulvic acid on the HSA hematite demonstrating the impact the second oxygen shell and two iron shells have on the overall shape of the EXAFS spectrum.
150 4 4 Citric acid Citric acid (001) (001) pH 4 pH 6 3 3
2 2
1 1
0 0 0 2 4 6 8 10 0 2 4 6 8 10 4 4 Citric acid Citric acid
3 (110) (110)
/Å pH 4 pH 6
− 3 3
2 2
1 1 Electron Density e
0 0 0 2 4 6 8 10 0 2 4 6 8 10 4 4 Citric acid Citric acid (012) (012) pH 4 pH 6 3 3
2 2
1 1
0 0 0 2 4 6 8 10 0 2 4 6 8 10 Height above Surface (Å) Height above Surface (Å)
Lead Total Lead only
Figure 4.7. XR (thick black lines) and RAXR results of lead and citric acid (red area) as well as the lead only RAXR from chapter 3 (gray line) on the three hematite surfaces.
151
LSA Pb (0.7) − Phthalic acid (7 mM) pH 6
LSA Pb (0.7) − Phthalic acid (7 mM) pH 4 ) 4 − (k)
HSA Pb (1) − Phthalic acid (10 mM) pH 6 3 k (R)| (Å |
HSA Pb (1) − Phthalic acid (10 mM) pH 4
2 3 4 5 6 7 8 9 10 0 1 2 3 4 5 −1 k (Å ) R (Å)
Figure 4.8. EXAFS of phthalate and lead adsorbed on hematite at listed pH. Data in blue, fit in red. The vertical line at k = 3.95 Å-1 is the antinode location of the lead only adsorbed on hematite and the vertical line at k = 3.45 Å-1 is the location of the Pb-Citrate antinode.
152 4 4 Phthalic acid Phthalic acid (001) (001) pH 4 pH 6 3 3
2 2
1 1
0 0 0 2 4 6 8 10 0 2 4 6 8 10 4 4 Phthalic acid Phthalic acid
3 (110) (110)
/Å pH 4 pH 6
− 3 3
2 2
1 1 Electron Density e
0 0 0 2 4 6 8 10 0 2 4 6 8 10 4 4 Phthalic acid Phthalic acid (012) (012) pH 4 pH 6 3 3
2 2
1 1
0 0 0 2 4 6 8 10 0 2 4 6 8 10 Height above Surface (Å) Height above Surface (Å)
Lead Total Lead only
Figure 4.9. XR (thick black lines) and RAXR results of lead and phthalic acid (red area) as well as the lead only RAXR from chapter 3 (gray line) on the three hematite surfaces.
153 LSA Pb (0.7) − Fulvic acid (60 mgC/L) pH 6
LSA Pb (0.7) − Fulvic acid (60 mgC/L) pH 4 ) 4 − (k)
HSA Pb (1) − Fulvic acid (85 mgC/L) pH 6 3 k (R)| (Å |
HSA Pb (1) − Fulvic acid (85 mgC/L) pH 4
2 3 4 5 6 7 8 9 10 0 1 2 3 4 5 −1 k (Å ) R (Å)
Figure 4.10. EXAFS of Fulvic acid and lead adsorbed on hematite at listed pH. Data in blue, fit in red. The vertical line at k = 3.95 Å-1 is the antinode location of the lead only adsorbed on hematite and the vertical line at k = 3.45 Å-1 is the location of the Pb-Citrate antinode.
154
4 4 Fulvic acid Fulvic acid (001) (001) pH 4 pH 6 3 3
2 2
1 1
0 0 0 2 4 6 8 10 0 2 4 6 8 10 4 4 Fulvic acid Fulvic acid
3 (110) (110)
/Å pH 4 pH 6
− 3 3
2 2
1 1 Electron Density e
0 0 0 2 4 6 8 10 0 2 4 6 8 10 4 4 Fulvic acid Fulvic acid (012) (012) pH 4 pH 6 3 3
2 2
1 1
0 0 0 2 4 6 8 10 0 2 4 6 8 10 Height above Surface (Å) Height above Surface (Å)
Lead Total Lead only
Figure 4.11. XR (thick black lines) and RAXR results of lead and fulvic acid (red area) as well as the lead only RAXR from chapter 3 (gray line) on the three hematite surfaces.
155
LSA Pb (0.7) − Humic acid (60 mgC/L) pH 6
LSA Pb (0.7) − Humic acid (60 mgC/L) pH 4 ) 4 − (k)
HSA Pb (1) − Humic acid (85 mgC/L) pH 6 3 k (R)| (Å |
HSA Pb (1) − Humic acid (85 mgC/L) pH 4
2 3 4 5 6 7 8 9 10 0 1 2 3 4 5 −1 k (Å ) R (Å)
Figure 4.12. EXAFS of humic acid and lead adsorbed on hematite at listed pH. Data in blue, fit in red. The vertical line at k = 3.95 Å-1 is the antinode location of the lead only adsorbed on hematite and the vertical line at k = 3.45 Å-1 is the location of the Pb-Citrate antinode.
156 4 (001) pH6 3
2
1
0 0 2 4 6 8 10 4
3 (110) pH6 /Å
− 3
2
1 Electron Density e
0 0 2 4 6 8 10 4 (012) pH6 3
2
1
0 0 2 4 6 8 10 Height above Surface (Å)
Lead Total Lead only Figure 4.13. XR (thick black lines) and RAXR results of lead and humic acid (red area) as well as the lead only RAXR from chapter 3 (gray line) on the three hematite surfaces.
157 Table 4.1. Results of the EXAFS model fitting
O C Fe CN R CN R CN R LSA Lead 2.58 2.31 (0) 0.64 3.32 (0.02) pH 6 0.83 3.92 (0.02) HSA Lead 2.49 2.30 (0.01) 0.76 3.34 (0.02) pH 6 0.75 3.94(0.02) LSA Citrate 2.19 2.33 (0.02) 3.85 3.23 (0.02) 0.63 3.44 (0.02) pH4 2.34 2.49 (0.02) 0.72 4.01 (0.03) LSA Citrate 2.3 2.29 (0.02) 1.75 3.2 (0.05) 0.54 3.37 (0.03) pH 6 1.99 2.47 (0.02) 0.77 3.97 (0.03) HSA Citrate 3.01 2.38 (0.01) 2.6 3.35 (0.02) 0.5 3.54 (0.02) pH3 2.41 2.58 (0.01) HSA Citrate 2.27 2.32 (0.01) 3.23 3.29 (0.02) 0.37 3.50 (0.04) pH4 2.16 2.50 (0.02) 0.37 4.00 (0.06) HSA Citrate 2.27 2.30 (0.01) 2.28 3.24 (0.03) 0.43 3.41 (0.04) pH 6 1.68 2.47 (0.02) 0.51 3.95 (0.04) LSA Phthalate 1.86 2.30 (0.01) 1.55 3.28 (0.03) 0.40 3.48 (0.02) pH4 2.47 2.46 (0.01) 1.02 3.99 (0.01) LSA Phthalate 2.62 2.33 (0.01) 0.93 3.21 (0.08) 0.65 3.39 (0.03) pH 6 0.88 2.50 (0.05) 0.69 3.95 (0.03) HSA Phthalate 2.51 2.36 (0.02) 1.48 3.27 (0.08) 0.51 3.43 (0.05) pH4 1.99 2.55 (0.03) 0.47 4.02 (0.07) HSA Phthalate 1.52 2.26 (0.03) 1.76 3.22 (0.03) 0.38 3.36 (0.03) pH 6 1.36 2.38 (0.03) 0.50 3.90 (0.03) LSA Fulvic 1.88 2.28 (0.02) 2.59 3.16 (0.02) 0.42 3.37 (0.02) pH4 1.67 2.43 (0.02) 0.67 3.95 (0.02) LSA Fulvic 2.54 2.31 (0.01) 1.38 3.04 (0.03) 0.78 3.28 (0.02) pH 6 1.01 3.9 (0.02) HSA Fulvic 1.29 2.23 (0.04) 3.15 3.17 (0.03) 0.61 3.37 (0.03) pH4 2.03 2.39 (0.03) 0.55 3.94 (0.05) HSA Fulvic 2.45 2.29 (0.00) 1.38 3.08 (0.03) 0.85 3.3 (0.02) pH 6 0.71 3.93 (0.03) LSA Humic 2.02 2.28 (0.02) 1.15 3.29 (0.05) 0.65 4.00 (0.03) pH4 2.30 2.46 (0.02) LSA Humic 2.81 2.3 (0.01) 0.40 3.3 (0.02) pH 6 0.98 2.49 (0.02) 0.61 3.9 (0.02) HSA Humic 2.51 2.32 (0.01) 3.06 3.26 (0.02) 0.37 3.45 (0.03) pH4 1.23 2.51 (0.02) HSA Humic 2.51 2.29 (0.01) 1.95 3.09 (0.03) 0.69 3.31 (0.02) pH 6 0.53 3.88 (0.04)
158
Table 4.2. Result from the linear combination fitting
LSA Adsorbed Lead Lead Lead Reduced pH Lead Citrate Fulvic Phthalate Pb2+ R-Factor χ2 χ2 4 21.1 64.8 -- -- 14.1 0.2909 9.9134 0.0909 Citrate 6 47.8 47.1 -- -- 5.2 0.2862 8.4238 0.0773 4 54.1 -- -- 26.9 19 0.4188 15.5401 0.1426 Phthalate 6 74.5 -- -- 10.7 14.7 0.2585 9.6644 0.0887 4 56.1 -- 39.4 -- 4.5 0.2468 7.4653 0.0685 Fulvic 6 82.4 -- 17.2 -- 0.4 0.1824 7.0492 0.0647 4 56.5 -- 33.4 0 10.2 0.3147 9.4873 0.0878 Humic 6 86.2 -- 8.5 0 5.3 0.1316 4.8554 0.0450 HSA 3 26.9 56.1 -- -- 17 0.2367 6.8945 0.0633 Citrate 4 33.6 56.1 -- -- 10.3 0.1975 5.2954 0.0486 6 54.9 32.3 -- -- 12.8 0.1389 3.4684 0.0318 4 50.1 -- -- 36.1 13.7 0.2930 7.8427 0.0720 Phthalate 6 78.5 -- -- 10 11.5 0.0999 2.7946 0.0256 4 63.3 -- 21.7 -- 15 0.1841 4.5911 0.0421 Fulvic 6 85.5 -- 10.7 -- 3.8 0.0424 1.3638 0.0125 4 63.9 -- 24.6 0 11.5 0.1524 4.1838 0.0384 Humic 6 87.1 -- 8.2 0 4.7 0.0523 1.8835 0.0173
159 Table 4.3. Results of Model dependent RAXR fit of lead on three faces of hematite in the presence of citric acid.
First peak Second peak Third peak Total χ2 R-Factor Crystal pH z (Å) c µ z (Å) c µ z (Å) c µ (µmol/m2) (001) 4 8.10 0.021 1.33(5) 1.20(7) 0.2(f) 3.72(12) 0.88(13) 0.79(19) 2.09(15) (001) 6 9.23 0.016 1.97(11) 0.48(24) 0.36(32) 4.26(21) 2.31(34) 1.49(19) 2.78(42) (110) 4 16.19 0.034 0.54(2) 2.10(5) 0.2(f) 3.22(6) 0.84(10) 0.2(f) 6.33(12) 2.48 (23) 1.85(15) 5.42(25) (110) 6 4.01 0.024 0.56(1) 2.53(4) 0.33(2) 3.04(2) 2.38(4) 0.30(2) 5.73(5) 0.70(4) 0.45(9) 5.61(7) (012) 4 5.18 0.021 1.41(6) 0.54(4) 0.2(f) 4.22(3) 0.95(5) 0.2(f) 7.63(3) 2.66(15) 1.01(5) 4.14(16) (012) 6 5.05 0.017 1.13(5) 1.74(9) 0.96(6) 3.95(2) 1.88(6) 0.31(3) 6.3(6) 0.44(4) 0.2(f) 4.06(12) z = height above the surface (Å); c = occupancy of the layer (µmol/m2); µ = rms width 160
160
Table 4.4. Results of Model dependent RAXR fit of lead on three faces of hematite in the presence of phthalic acid.
First peak Second peak Third peak Total χ2 R-Factor Crystal pH z (Å) c µ z (Å) c µ z (Å) c µ (µmol/m2) (001) 4 2.32 0.006 1.61(3) 0.34(2) 0.2(f) 4.95(15) 0.52(10) 2.00(26) 0.86(10) (001) 6 5.38 0.013 2.15(5) 0.47(7) 0.34(13) 5.54(43) 0.64(18) 2.34(51) 1.11(19) (110) 4 7.26 0.033 0.83(3) 1.16(6) 0.27(5) 3.29(4) 1.51(9) 0.73(9) 6.76(14) 0.25(5) 0.32(26) 2.92(12) (110) 6 12.42 0.052 0.74(3) 1.24(6) 0.2(f) 3.03(4) 2.52(25) 0.59(6) 6.22(20) 3.15(40) 1.87(21) 6.91(47) (012) 4 3.64 0.021 0.92(5) 0.56(3) 0.2(f) 4.29(7) 1.06(7) 1.38(10) 1.62(7) (012) 6 2.94 0.017 1.62(6) 0.45(4) 0.2(f) 4.00(2) 1.4(6) 0.2(0) 6.41(8) 3.17(17) 1.94(9) 5.02(18) z = height above the surface (Å); c = occupancy of the layer (µmol/m2); µ = rms width 161
161
Table 4.5. Results of Model dependent RAXR fit of lead on three faces of hematite in the presence of fulvic acid. At pH 4, the (012) surface required an additional peak to fit the data.
First peak Second peak χ2 R-Factor Crystal pH z (Å) c µ z (Å) c µ (001) 4 1.73 0.007 2.95(7) 1.24(14) 1.65(10) 7.75(19) 1.94(4) 2.29(33) (001) 6 3.26 0.013 2.01(3) 1.88(8) 0.73(3) 5.35(5) 1.33(13) 0.94(8) (110) 4 11.67 0.041 0.65(4) 1.69(7) 0.42(6) 2.94(3) 1.85(9) 0.35(6) (110) 6 9.90 0.040 0.25(7) 0.59(5) 0.2(f) 3.36(4) 1.35(10) 0.26(8) (012) 4 1.89 0.012 2.28(8) 0.56(4) 0.45(f) 5.00(5) 0.59(3) 0.2(f) (012) 6 8.15 0.021 1.78(2) 1.75(4) 0.33(f) 4.57(6) 0.55(5) 0.2(f)
Third peak Diffuse Total Crystal pH z (Å) c µ z (Å) c µ (µmol/m2) (001) 4 3.18(32) (001) 6 15.25(59) 0.36(2) 1(f) 3.57(22) 162 (110) 4 5.62(9) 2.42(5) 1(f) 5.96(23)
(110) 6 5.00(7) 5.87(4) 1(f) 7.8(20) (012) 4 8.83(8) 0.62(5) 0.54 (f) 11.93(46) 3.07(4) 2.85(f) 4.84(50) (012) 6 6.02(8) 3.28(1) 1(f) 5.57(13) z = height above the surface (Å); c = occupancy of the layer (µmol/m2); µ = rms width
162
Table 4.6. Results of Model dependent RAXR fit of lead on three faces of hematite in the presence of Humic acid.
First peak Second Peak Third Peak Total Crystal pH χ2 R-Factor z (Å) c µ z (Å) c µ z (Å) c µ (µmol/m2) (001) 6 1.97 0.008 1.97(5) 0.46(3) 0.2(f) 5.32(19) 1.64(4) 2.35(42) 2.09(28) (110) 6 9.34 0.041 0.86(5) 0.97(9) 0.2(f) 3.54(7) 1.71(16) 0.78(16) 6(23) 0.31(11) 0.33(25) 2.98(21) (012) 6 2.75 0.019 1.27(3) 1.85(7) 0.75(6) 4.15(5) 1.35(32) 0.39(11) 6.21(33) 1.92(38) 1.39(24) 5.12(50) z = height above the surface (Å); c = occupancy of the layer (µmol/m2); µ = rms width
163
163
Chapter 5: Conclusions and Future Work
The following are the objectives set at the beginning of the project and the progress made on meeting those objectives. Following that are ideas for future research to further the knowledge of organic acid and heavy metal interactions at the mineral – water interface.
5.1 Objectives
1. Determine the bonding mode of citric acid on hematite nanoparticles and the impact of particle size.
In Chapter 2, the combinations of infrared spectroscopy, batch adsorption, density functional theory molecular modeling, and surface complexation modeling were used to investigate the adsorption of citrate onto the hematite surface. Citrate adsorption on goethite has been investigated in several recent studies, however a consensus has not been reached. Hematite has not been used as a sorbent for citrate in a study since 1985
(see Table 1.1). Our results suggest that citrate adsorbs as both inner-sphere and outer- sphere complex at low pH and an outer-sphere complex under higher pH conditions. The inner-sphere complex is completely deprotonated, including the hydroxyl group which has been implicated in iron – citrate complexes and in citrate adsorption to goethite. The outer-sphere complex is singly protonated at low pH and deprotonated at higher pH.
These three adsorption modes were included in a triple layer surface complexation model
164 to determine the equilibrium coefficients. Once determined, the coefficients were tested for a variety of solution conditions and particle sizes and returned adequate results, though on the higher pH side of the adsorption envelop, the model did not fit the experimental results in those other conditions perfectly indicating there is an additional adsorption mode we did not account for. The size of the particles had a minimal impact on the adsorption envelop of citrate on hematite when normalized to surface area. The
FTIR and modeling indicated a slight preference for the inner-sphere complex on the smaller hematite particles as would be expected based on the size-dependent changes in capacitance.
2. Further the knowledge of lead adsorption on hematite with single crystal studies.
Previous studies concluded that lead adsorbs onto the hematite surface as an inner-sphere complex as both edge sharing and corner sharing bidentate structures. The particle based
(extended X-ray adsorption fine structure – EXAFS) results in chapter 3 confirm this.
Lead adsorption on the particle surface increases sharply with pH between approximately
4 and 6 depending on the solution conditions, and the single crystal resonant anomalous
X-ray reflectivity (RAXR) experiments were thus carried out at pH 4 and 6. The single crystal experiments show where the lead was adsorbed as a function of lateral distance from each surface. The results from the RAXR experiments indicate in addition to the inner-sphere complex, that there was also an outer-sphere complex that forms on the
(012) and (110) at pH 6 and (012) at pH 4 that accounts for a nontrivial amount of lead adsorbed. The (001) surface at both pH conditions has a low level of lead adsorbed in a broad distribution. At pH 6, lead adsorbs on the different surfaces in the order (012) >
165 (110) > (001). At pH 4, this is switch to (110) > (012) > (001). Adsorption on the (012) surface was the most variable with pH while the (001) did not vary. The (012) and (110) have a corrugated topography with singly and triply coordinated oxygen atoms while the
(001) is atomically flat with only doubly coordinated oxygen. The corrugated topography and the singly and triply coordinated oxygen atoms that were present on the
(012) and (110) surfaces were likely the reason for the increase in reactivity. This was also relevant to the (001) surface and any reactivity found there was likely due to surface defects disrupting the flat surface.
3. Determine the effect of organic acids on the adsorption of lead to hematite on both nanoparticles and specific crystal faces.
Chapter 4 used the same methods as chapter 3 to investigate the impact that the organic acids citric acid, phthalic acid, humic acid and fulvic acid have on the adsorption of lead to hematite. The presence of organic acids has been reported to shift the adsorption edge to lower pH values. This was seen with citric acid in batch adsorption experiments on both the LSA and HSA hematite particles. At higher lead and citrate concentrations, the citrate inhibits the adsorption of lead at pH 6. The other three acids have minimal effect at pH 6 and increase adsorption at pH 4. The general results of the RAXR was that the surface plays as bigger role than the organic acids do in the adsorption process. The
EXAFS and RAXR both showed that the inner sphere complexation continued to play a significant, if not dominant, role in the lead adsorption. The acids, especially the fulvic acid increased the amount of lead adsorbed at a distance from the surface. These
166 complexes were either electrostatically held as a result of the acid decreasing the surface charge or held in a ternary complex.
4. Merge the results of the single crystal work with the particle based studies.
The combination of particle and surface based techniques yields a clearer view of the adsorption process than either one would have independently. The EXAFS helped inform and set limits on the possible adsorption modes derived from RAXR that only provide distance of the lead from the hematite surface. The XR / RAXR also give information on the amount of lead that was adsorbed in the different modes that was not available from the EXAFS. Outer-sphere complexes, which play a significant role, especially when acids are added, were directly visible in the single crystal studies whereas they need to be inferred in the EXAFS.
5.2 Future Work
1. There are many directions the work can go on from here. The use of theoretical methods, either ab initio molecular modeling, to determine the most likely position of the adsorbates on hematite, or molecular dynamics simulations to determine the changes of the adsorbates with times would be a natural compliment to the experimental work on single crystals. One of the advantages of the single crystal methods is the more precise knowledge of surface composition. This will allow for improved results over the small cluster model used in Chapter2.
2. There was limited research done into the effect of particle size on the hematite surface using a 99 m2/g hematite and a 35 m2/g hematite. The biggest difference between the two particle sizes was the shift to lower pH values in the adsorption edge of lead on the 167 smaller particles. Both of these particles were defined as roughly spherical. Hematite can be synthesized in many different shapes and sizes. Performing adsorption, IR and
EXAFS experiments on a wider variety of hematite particles may yield more variable results than we found. Studying particles with more clearly defined crystal surfaces, such as the platy crystals with large fraction of (001) faces, may be the most interesting to compare to the more spherical particles we studied. In addition to different shaped particles, changing the solution conditions may yield more information on the system.
Almost all of the studies were performed in a 0.1M background electrolyte concentration and at concentrations of adsorbates intended to saturate the surface. The limited experiments with varying the concentration of citrate (chapter 2) showed that citrate has multiple adsorption modes on hematite.
3. Surface complexation modeling was performed in chapter 2 with citrate on the hematite surface. The next step would be to include lead and new information about the hematite site types in the model. Taking what was learned from the single crystal studies, and including a more complex surface, the reactivity of which is dominated by singly and triply coordinated oxygen atoms, may improve the result of the surface complexation model.
4. One issue touched on, but not explicitly explored, in chapter 4 was the affect of the order of addition on the ternary systems. Components of the EXAFS and batch experiment were added as close to each other as possible. Under these co-addition conditions the lead inner-sphere complex that formed in the absence of the organic acid was still a prominent adsorption mode. The organic acids tested have different
168 adsorption tendencies. Citric acid adsorbs mostly as an outer-sphere complex where as phthalic acid is more an inner-sphere adsorption. Both acids have been known to chelate lead in solution, so allowing the hematite to equilibrate with either lead or acid first, before adding the other may yield different results, shedding more light on the hematite – lead – organic acid ternary system
169
Appendix A: Supporting Information for Chapter 2
(A) (B)
Figure A.1. TEM images of (A) high surface area (99 m2/g) and (B) low surface area (35 m2/g) hematite
170
A B
C D
E F
G H
Figure A.2. Optimized structures for A) protonated and B) deprotonated outer-sphere complex (OS-BN), C) protonated and D) deprotonated mononuclear bidentate complex bound by the central and one terminal carboxyl (CT-MN), E) protonated and F) deprotonated binuclear bidentate complex with the central and terminal carboxyl groups (CT-BN), and G) protonated and H) deprotonated binuclear bidentate complex bound by the central carboxyl and deprotonated hydroxyl group (CH-BN). Atoms outlined in red are part of the iron oxide cluster, atoms outlined in blue are the explicit water molecules and the atoms outlined in green are the citrate molecule.
171
Explicit Water
IEFPCM only
Experimental
1800 1700 1600 1500 1400 1300 1200 −1 Wavenumber (cm ) Figure A.3. Comparison of theoretical citrate spectra with and without (IEFPCM only) explicit water molecules. Without explicit water molecules, the C-O-H bending peak at1468 cm-1 is too intense and at too high of a wavenumber. Adding the explicit water molecules results in a more realistic spectrum.
172 1800 CH−BN CT−BN 1750 CT−MN OS−BN ) 1
− 1700
1650 Carb 1600
1550 Asym
1500
1450
1400 Theoretical Wavenumber (cm 1350 Sym 1300 1800 1750 1700 1650 1600 1550 1500 1450 1400 1350 1300 −1 Experimental Wavenumber (cm ) Figure A.4. Comparison of the theoretical asymmetric and symmetric stretch peak locations with the experimental locations on the low surface area hematite at pH 2.5 and 7.5
173 (A) Fe
O H H O O O Fe H C H H C H O O H H C H Fe O C H H O C O O H H C O Fe H O O H
Fe
(B) Fe
O O O O C Fe Fe O C O C H H O C Fe Fe
C H O C H O Fe O
O H
Fe
Figure A.5. Simplified representation of (A) Outer-sphere and (B) inner-sphere complexes of citrate on hematite. The Inner-sphere complex does no occur on the atomically flat (001) hematite surface. The additional iron layer on the inner-sphere complex represents the more corrugated structure of other common hematite faces.
174
Table A.1. Solution species chemistry used for the surface complexation modeling. eq Formula Log K Ref - + 1 H3Citrate --> H2Citrate + H -3.13 (Stumm and Morgan, 1996) 2- + 2 H3Citrate --> HCitrate + 2H -7.89 (Stumm and Morgan, 1996) 3- + 3 H3Citrate --> Citrate + 3H -14.29 (Stumm and Morgan, 1996) 4- + 4 H3Citrate --> H-1Citrate + 4H -28.69 (Silva et al., 2009) + 2- + 5 H3Citrate + Na --> NaCitrate + 3H -12.79 (Stumm and Morgan, 1996)
175 Table A.2. Coordinates for protonated outer-sphere complex shown in figure A.2. A Atom X Y Z Atom X Y Z Fe -1.799089 -1.415614 0.853496 ! H 3.628883 3.573336 -0.607308 Fe -4.083537 -0.352185 -0.634194 ! C 4.595933 1.127858 -1.300697 O -3.322091 -0.443745 1.024513 ! H 5.312751 1.945209 -1.193683 H -3.515632 0.158864 1.766522 ! H 4.309153 1.052151 -2.354753 O -2.476340 -1.403824 -0.990134 ! C 5.258961 -0.176418 -0.909779 H -2.661931 -2.286859 -1.351481 ! O 6.544876 -0.041495 -0.605917 O -5.262673 -1.734966 -0.463147 ! O 4.661869 -1.252029 -0.899478 H -4.841803 -2.589265 -0.268877 ! C 1.578568 3.266472 -0.033847 O -5.603781 0.809035 -0.011342 ! O 0.574241 3.700913 -0.664451 H -6.384754 0.229239 0.071245 ! O 1.642215 3.117811 1.232858 O -4.671118 -0.362342 -2.570464 ! C 3.341991 1.441979 -0.460897 H -5.469531 -0.919899 -2.627829 ! O 3.735128 1.409256 0.904708 O -2.33376 -3.092795 1.215940 ! H 3.051408 1.931942 1.385848 H -2.393046 -3.670524 0.434489 ! C 2.199003 0.424055 -0.771192 O -1.261722 -1.327907 2.780435 ! O 1.690645 -0.220017 0.221984 H -1.344832 -2.224163 3.153394 ! O 1.784365 0.352842 -1.946181 O -0.066048 -2.196471 0.187010 ! H 6.981049 -0.922837 -0.378123 H 0.256680 -2.919966 0.751374 ! O 5.588148 -3.818212 -0.308083 H -4.887628 0.501366 -2.964525 ! H 5.126511 -2.968049 -0.494486 H -5.874706 1.567054 -0.558886 ! H 5.175621 -4.179204 0.492878 H -0.305476 -1.051113 2.893162 ! O 7.820469 -2.305789 0.067294 H 0.688383 -1.532554 0.092721 ! H 7.143351 -3.024105 -0.027891 O -3.102154 1.301553 -1.282110 ! H 8.535157 -2.520105 -0.553929 H -2.810321 1.967513 -0.610094 ! O -0.550361 3.78281 2.736193 H -2.272741 1.059693 -1.785318 ! H 0.289311 3.558645 2.246479 O -0.825320 0.340242 0.794395 ! H -0.551757 3.227442 3.531597 H -1.211060 1.244472 0.635631 ! O 1.347869 -0.544138 2.958208 H 0.118104 0.331790 0.499429 ! H 1.706457 -0.362137 2.057023 O -1.725512 2.866131 0.441923 ! H 1.950991 -1.187462 3.363395 H -1.612408 3.187848 1.370757 ! H -0.923985 3.250819 -0.023542 ! ! ! ! ! O -0.979105 0.217875 -2.620473 ! ! ! ! ! H -1.320183 -0.574199 -2.145588 ! ! ! ! ! H -0.029193 0.287897 -2.364469 ! ! ! ! ! C 2.825481 2.861226 -0.834321 ! ! ! ! ! H 2.607105 2.911007 -1.902067 ! ! ! ! ! ! ! ! ! !
176 Table A.3. Coordinates for deprotonated outer-sphere complex shown in figure A.2 B. Atom X Y Z Atom X Y Z Fe -2.204834 -1.259938 1.096982 ! H 2.583145 2.538611 -1.791939 Fe -4.130758 -0.078394 -0.724740 ! H 3.723827 3.178313 -0.594449 O -3.636860 -0.136276 1.034882 ! C 4.398622 0.670012 -1.055620 H -4.013702 0.390994 1.761852 ! H 5.185090 1.420291 -0.918750 O -2.475317 -1.103646 -0.837562 ! H 4.150320 0.671809 -2.124049 H -2.652659 -1.984665 -1.211043 ! C 5.019924 -0.704724 -0.759609 O -5.207045 -1.542943 -0.639553 ! O 6.191776 -0.875172 -1.206819 H -4.799076 -2.219734 -0.064630 ! O 4.338706 -1.582768 -0.142826 O -5.727059 1.091392 -0.379900 ! C 1.651900 3.235172 0.013232 H -6.557560 0.592105 -0.485345 ! O 0.504099 2.730400 -0.197720 O -4.489850 -0.176235 -2.706292 ! O 1.874224 4.236485 0.750586 H -5.155597 -0.880323 -2.824880 ! C 3.156228 1.130955 -0.254257 O -3.060038 -2.859441 1.149322 ! O 3.465054 1.251090 1.135150 H -2.689111 -3.483253 0.499606 ! H 3.473976 0.367851 1.557182 O -2.214868 -1.333593 3.106143 ! C 1.943696 0.188958 -0.536122 H -2.310300 -2.247796 3.429662 ! O 1.349792 -0.357768 0.470801 O -0.493842 -2.235329 0.720318 ! O 1.615489 0.023599 -1.729105 H -0.169548 -2.798375 1.444379 ! O 5.191921 -4.258515 -0.122391 H -4.831099 0.618031 -3.154473 ! H 4.843425 -3.334470 -0.110192 H -5.815060 1.904265 -0.908326 ! H 6.094064 -4.132530 -0.487901 H -1.409500 -0.960626 3.507899 ! O 7.602372 -3.186102 -1.166260 H 0.263496 -1.602927 0.484958 ! H 7.086766 -2.329766 -1.190642 O -3.082166 1.586930 -1.220331 ! H 7.805926 -3.396964 -2.090878 H -2.743465 2.119673 -0.451418 ! O 4.248125 5.452215 1.321994 H -2.263204 1.338393 -1.739310 ! H 3.416173 4.957262 1.094548 O -1.015019 0.319813 1.474410 ! H 4.874965 4.776137 1.622062 H -1.392935 1.228779 1.333567 ! O 3.183961 -1.366585 2.346162 H -0.115292 0.302510 1.049483 ! H 2.327546 -1.132882 1.931091 O -1.820854 2.813930 0.831777 ! H 3.710244 -1.633265 1.558023 H -2.085630 3.507672 1.456175 ! H -0.876067 3.011227 0.492877 ! ! ! ! ! O -1.019910 0.426846 -2.559044 ! ! ! ! ! H -1.397664 -0.329178 -2.053819 ! ! ! ! ! H -0.077473 0.440989 -2.266706 ! ! ! ! ! C 2.822939 2.571604 -0.724764 ! ! ! ! ! ! ! ! ! !
177
Table A.4. Coordinates for protonated inner-sphere mononuclear bidentate complex (CT-MN) shown in figure A.2 C
Atom X Y Z Atom X Y Z Fe -1.541283 -1.997868 0.126175 ! O 5.176976 0.068840 1.637939 O -3.417407 -2.450480 0.516124 ! H 4.446226 0.521352 2.138367 H -3.908701 -3.046990 -0.079584 ! H 5.533816 -0.618644 2.223604 O -1.427940 -1.416547 2.003657 ! O -0.518798 3.424090 0.619496 O -1.766994 -2.735369 -1.692530 ! H -1.338916 3.104917 0.191753 H -2.342323 -2.249583 -2.313230 ! H -0.673425 3.350235 1.575050 O -1.013757 -3.832871 0.697116 ! O 2.967524 1.324434 2.649099 H -1.332447 -4.579808 0.156046 ! H 2.757429 1.212168 3.590108 H -1.395133 -2.168605 2.625040 ! H 2.203146 0.951013 2.161963 H -0.936214 -2.921876 -2.170311 ! O -4.831405 1.977304 0.527799 H -0.044156 -3.933594 0.758921 ! H -3.979772 1.909362 0.048977 C -0.439547 0.635467 -1.793205 ! H -5.142423 2.896715 0.391242 H -0.534742 -0.253246 -2.421555 ! O -5.869513 4.626089 0.213072 H -0.383407 1.509500 -2.444870 ! H -5.317814 5.326876 0.597065 C 2.007783 0.020329 -1.979179 ! H -6.012299 4.895882 -0.708775 H 2.017335 0.643535 -2.875451 ! H 1.778603 -1.010567 -2.264330 ! ! ! ! ! C 3.379265 0.071180 -1.336365 ! ! ! ! ! O 4.307279 0.750130 -1.753066 ! ! ! ! ! O 3.442947 -0.707543 -0.254796 ! ! ! ! ! C -1.704842 0.764217 -0.943235 ! ! ! ! ! O -2.151256 -0.255521 -0.254734 ! ! ! ! ! O -2.325247 1.838111 -0.912317 ! ! ! ! ! C 0.921129 0.552897 -1.013064 ! ! ! ! ! O 1.373736 1.804561 -0.567795 ! ! ! ! ! H 0.689607 2.295308 -0.047197 ! ! ! ! ! C 0.679484 -0.412734 0.162165 ! ! ! ! ! O 0.777169 -0.073201 1.342566 ! ! ! ! ! O 0.219073 -1.594434 -0.216756 ! ! ! ! ! H -3.960110 -1.644382 0.620796 ! ! ! ! ! H -0.572674 -0.916860 2.108889 ! ! ! ! ! H 4.248221 -0.518107 0.311571 ! ! ! ! ! O 6.668831 1.609066 -0.265729 ! ! ! ! ! H 6.359880 1.170294 0.551702 ! ! ! ! ! H 5.964270 1.378276 -0.902202 ! ! ! ! ! ! ! ! ! !
178
Table A.5. Coordinates for protonated inner-sphere mononuclear bidentate complex (CT-MN) shown in figure A2 D.
Atom X Y Z Atom X Y Z Fe 1.160633 2.341458 -0.026990 ! O -5.644416 0.308697 0.808787 O 2.780325 3.445977 0.214814 ! H -4.715706 0.049159 0.603711 H 3.106873 3.989240 -0.526847 ! H -6.143626 -0.118589 0.079614 O 1.084231 2.256611 1.933860 ! O -2.557897 -2.706217 1.720646 O 1.306834 2.635358 -1.981824 ! H -2.985879 -1.922165 1.312242 H 2.019438 2.170395 -2.458752 ! H -2.197977 -2.389654 2.580584 O 0.055294 4.009977 0.058771 ! O -1.137999 -1.529577 3.829318 H 0.210558 4.708762 -0.603919 ! H -0.439043 -2.107341 4.175477 H 0.681524 3.060677 2.313794 ! H -0.714414 -0.986398 3.133454 H 0.485529 2.477702 -2.484623 ! O 3.791313 -3.866761 -0.831109 H -0.899701 3.808453 0.024505 ! H 3.544053 -2.951611 -0.584496 C 1.069350 -0.859930 -1.240947 ! H 4.745895 -3.947221 -0.623972 H 0.965329 -0.155517 -2.068566 ! O 6.577313 -4.228593 -0.264161 H 1.320068 -1.836538 -1.658951 ! H 7.164471 -3.633196 -0.757837 C -1.414849 -1.027096 -1.623803 ! H 6.809225 -4.092631 0.668979 H -1.268170 -1.917469 -2.243431 ! H -1.292767 -0.158214 -2.281458 ! ! ! ! ! C -2.864694 -0.983272 -1.118144 ! ! ! ! ! O -3.754903 -1.370309 -1.921526 ! ! ! ! ! O -3.056650 -0.523860 0.053462 ! ! ! ! ! C 2.239159 -0.406862 -0.366786 ! ! ! ! ! O 2.289202 0.836537 0.069842 ! ! ! ! ! O 3.160476 -1.178450 -0.087841 ! ! ! ! ! C -0.319828 -1.013277 -0.535596 ! ! ! ! ! O -0.263572 -2.220211 0.192242 ! ! ! ! ! H -1.099848 -2.377701 0.699622 ! ! ! ! ! C -0.443685 0.198414 0.399007 ! ! ! ! ! O -0.417078 0.123850 1.630333 ! ! ! ! ! O -0.377331 1.360247 -0.244043 ! ! ! ! ! H 3.531342 2.899793 0.518623 ! ! ! ! ! H 0.467437 1.497324 2.142401 ! ! ! ! ! O -6.436118 -1.163868 -1.486897 ! ! ! ! ! H -6.797669 -0.661864 -2.234015 ! ! ! ! ! H -5.458308 -1.246430 -1.664205 ! ! ! ! ! ! ! ! ! !
179
Table A.6. Coordinates for protonated inner-sphere binuclear bidentate complex with the terminal and central carboxyl groups bound to the iron oxide cluster (CT-BN) shown in figure A.2 E. Atom X Y Z Atom X Y Z Fe 3.086297 -0.248280 0.238173 ! C -1.663384 -1.296046 -0.731563 Fe 1.105222 1.892997 0.372998 ! O -2.249759 -2.189565 0.205200 O 4.383167 0.489751 1.248135 ! H -1.517784 -2.560337 0.735633 H 5.129373 0.762470 0.684550 ! C -1.235627 0.022512 -0.006676 O 2.120134 3.154848 1.238018 ! O -1.787846 0.320918 1.076505 H 2.740918 3.578864 0.620566 ! O -0.310826 0.711348 -0.568229 O 1.735754 0.417732 1.469095 ! O -7.598977 -0.494967 0.080134 O 2.408107 1.160823 -0.707078 ! H -7.425321 0.464282 0.122660 O 3.183232 -1.901511 1.411730 ! H -6.763349 -0.839812 -0.293146 H 2.255164 -2.259164 1.322852 ! O -6.278204 2.081319 -0.167554 O 4.509335 -0.722538 -1.132105 ! H -6.072526 2.597503 0.630737 H 4.189358 -1.230478 -1.932983 ! H -6.641226 2.721903 -0.802860 O 0.516457 3.305413 -0.970282 ! O -2.819476 -1.200271 3.152472 H -0.434462 3.280407 -1.178812 ! H -3.443763 -1.887287 2.869821 O -0.287615 2.332549 1.770153 ! H -2.573591 -0.726575 2.323783 H -0.670556 3.223368 1.706522 ! O -0.321531 -2.473733 3.587000 H 3.769788 -2.601459 1.075294 ! H -1.198991 -2.034189 3.506544 H 5.278362 -1.207132 -0.787620 ! H 0.064888 -2.416806 2.695209 H 2.248617 0.829839 2.187007 ! O 1.076767 0.906136 -3.054021 H 2.211043 1.086574 -1.678884 ! H 0.362565 0.798744 -2.390684 H 0.699586 4.204450 -0.642039 ! H 0.894073 1.749818 -3.498702 H -1.026651 1.669594 1.632899 ! O 3.282346 -1.949032 -3.169633 C -0.430962 -1.966658 -1.396625 ! H 2.491148 -1.892965 -2.590324 H -0.158346 -1.459636 -2.322483 ! H 3.466813 -2.896244 -3.278702 H -0.712423 -2.997762 -1.637090 ! C -2.726630 -1.014434 -1.822228 ! ! ! ! ! H -2.942537 -1.961045 -2.323857 ! ! ! ! ! H -2.312591 -0.312451 -2.550910 ! ! ! ! ! C -4.034233 -0.459834 -1.290774 ! ! ! ! ! O -5.005601 -1.154794 -1.021087 ! ! ! ! ! O -4.014231 0.869021 -1.171325 ! ! ! ! ! H -4.869524 1.224829 -0.803257 ! ! ! ! ! C 0.757896 -1.962351 -0.445197 ! ! ! ! ! O 1.830621 -1.405544 -0.873823 ! ! ! ! ! O 0.625657 -2.459285 0.704971 ! ! ! ! ! ! ! ! ! !
180
Table A.7. Coordinates for deprotonated inner-sphere binuclear bidentate complex with the terminal and central carboxyl groups bound to the iron oxide cluster (CT-BN) shown in figure A.2 F.
Atom X Y Z Atom X Y Z Fe 3.086413 -0.081829 -0.054990 ! C -1.729068 -1.403375 -0.399301 Fe 1.009279 1.878436 0.528370 ! O -2.540698 -2.241426 0.401164 O 4.507950 0.744779 0.696905 ! H -2.002819 -2.735013 1.068325 H 5.112646 1.059315 0.001404 ! C -1.192587 -0.189262 0.414511 O 2.036203 3.201120 1.275126 ! O -1.445729 -0.089796 1.638450 H 2.433207 3.761520 0.585924 ! O -0.437662 0.628638 -0.229766 O 1.967452 0.458924 1.441130 ! O -7.463743 0.270525 -0.961767 O 2.150791 1.305031 -0.797832 ! H -7.881581 0.484741 -1.810452 O 3.560610 -1.736230 1.012485 ! H -6.558515 -0.093902 -1.188311 H 2.673415 -2.204507 1.072897 ! O -5.981565 2.347080 0.296695 O 4.230904 -0.433578 -1.699709 ! H -6.692340 1.784569 -0.078650 H 3.777741 -0.938368 -2.435088 ! H -5.166695 1.837948 0.054781 O 0.051716 3.261472 -0.625812 ! O -1.714956 -1.587887 3.965143 H -0.899553 3.074656 -0.720169 ! H -2.633039 -1.672122 4.267946 O -0.159022 2.102334 2.160171 ! H -1.749561 -1.025897 3.157940 H -0.710288 2.902655 2.161377 ! O -1.033201 -3.683556 2.249348 H 4.161511 -2.354668 0.562412 ! H -1.186134 -3.073070 3.007046 H 5.077069 -0.883353 -1.536904 ! H -0.186241 -3.394947 1.839800 H 2.580746 0.920809 2.040539 ! O 0.466127 0.931445 -2.875500 H 1.787572 1.243356 -1.721307 ! H -0.077399 0.734435 -2.079692 H 0.118160 4.167487 -0.274607 ! H 0.090148 1.741442 -3.256987 H -0.777212 1.309357 2.121324 ! O 2.659967 -1.660795 -3.486564 C -0.513525 -2.210980 -1.014661 ! H 1.999299 -1.629331 -2.759548 H -0.387579 -1.920165 -2.059084 ! H 2.845208 -2.602511 -3.634852 H -0.802204 -3.265815 -0.989365 ! C -2.630515 -0.947153 -1.568185 ! ! ! ! ! H -2.884081 -1.836188 -2.153076 ! ! ! ! ! H -2.067792 -0.269083 -2.217239 ! ! ! ! ! C -3.919306 -0.228077 -1.133616 ! ! ! ! ! O -5.018000 -0.729126 -1.524644 ! ! ! ! ! O -3.781743 0.825495 -0.448118 ! ! ! ! ! C 0.831767 -2.034698 -0.317736 ! ! ! ! ! O 1.679111 -1.278754 -0.919604 ! ! ! ! ! O 1.070205 -2.604926 0.780155 ! ! ! ! ! ! ! ! ! !
181
Table A.8. Coordinates for protonated inner-sphere binuclear bidentate complex with the central carboxyl group and deprotonated hydroxyl bound to the iron oxide cluster (CH- BN) shown in figure A.2 G.
Atom X Y Z Atom X Y Z Fe 0.958838 -1.265522 1.354998 ! O -0.148938 -0.098669 -1.515160 Fe 0.228322 -1.940221 -0.906501 ! C -0.517940 1.300715 0.457451 O -0.478392 -2.341417 0.807862 ! O -0.066777 0.358550 1.170876 H -0.197179 -3.256908 1.005608 ! O -1.047375 2.313276 0.993249 O 1.940370 -1.666042 -0.212315 ! H -4.845614 1.496959 -0.461310 H 2.290056 -2.570052 -0.078871 ! O 4.956455 2.757870 0.766707 O 0.622459 -3.736122 -1.116223 ! H 4.008629 2.553696 0.589892 H -0.185511 -4.264802 -0.994287 ! H 5.404923 2.307544 0.018541 O -1.667569 -2.299130 -1.592278 ! O 5.585365 1.361360 -1.626675 H -1.714465 -2.369233 -2.560983 ! H 4.598103 1.382109 -1.771785 O 1.058043 -1.727286 -2.807081 ! H 5.814267 0.428026 -1.495211 H 0.706383 -2.328980 -3.484536 ! O -5.273166 -1.502194 0.547356 O 1.749517 -2.750008 2.150991 ! H -5.649688 -2.218807 0.011663 H 2.718951 -2.765195 2.077586 ! H -4.484599 -1.183342 0.055894 O 0.140366 -1.034561 3.222632 ! O -6.296142 0.983969 0.157672 H 0.789165 -1.374752 3.861474 ! H -6.516548 1.445721 0.982989 O 2.441676 0.014225 1.848487 ! H -6.120161 0.039875 0.404668 H 2.390777 0.769115 1.183774 ! O -1.680059 4.917048 0.252807 H 0.660302 -0.829871 -2.932016 ! H -1.446025 3.984710 0.471025 H -2.162418 -1.465361 -1.368515 ! H -0.873893 5.427603 0.424916 H -0.060083 -0.083913 3.443502 ! O -0.730291 1.553189 3.749873 H 3.320780 -0.377958 1.712480 ! H -0.826801 1.874667 2.818781 C 0.628268 2.186091 -1.608021 ! H -0.121458 2.173353 4.181424 H 0.370099 3.198582 -1.276315 ! H 0.585504 2.159204 -2.701027 ! ! ! ! ! C -1.811704 1.648767 -1.688344 ! ! ! ! ! H -1.789609 1.357778 -2.745698 ! ! ! ! ! H -1.914265 2.733747 -1.640887 ! ! ! ! ! C -3.047174 1.036009 -1.075774 ! ! ! ! ! O -3.174989 -0.170461 -0.818584 ! ! ! ! ! O -4.008696 1.913174 -0.850420 ! ! ! ! ! C 2.062203 1.892335 -1.155919 ! ! ! ! ! O 2.313872 2.010804 0.094782 ! ! ! ! ! O 2.911463 1.597911 -2.038173 ! ! ! ! ! C -0.445868 1.179735 -1.076859 ! ! ! ! ! ! ! ! ! ! 182
Table A.9. Coordinates for deprotonated inner-sphere binuclear bidentate complex with the central carboxyl group and deprotonated hydroxyl bound to the iron oxide cluster (CH-BN) shown in figure A.2 H.
Atom X Y Z Atom X Y Z Fe 0.924298 -1.346791 1.323106 ! C -0.394008 1.195396 -1.034593 Fe 0.134671 -1.937002 -0.956848 ! O 0.053036 -0.045834 -1.470176 O -0.569776 -2.338764 0.765669 ! C -0.483328 1.283212 0.496859 H -0.339736 -3.275735 0.922321 ! O -0.046711 0.322408 1.196509 O 1.868435 -1.818000 -0.257585 ! O -0.996182 2.294650 1.053946 H 2.141595 -2.744157 -0.103493 ! O 5.004861 2.654485 0.854305 O 0.350823 -3.754684 -1.223106 ! H 4.055320 2.466546 0.664435 H -0.512202 -4.192625 -1.119749 ! H 5.449208 2.269465 0.068310 O -1.761649 -2.097691 -1.623973 ! O 5.609202 1.467352 -1.655601 H -1.820677 -2.023001 -2.590998 ! H 4.619957 1.513426 -1.787699 O 0.978274 -1.746138 -2.868370 ! H 5.825699 0.523377 -1.602808 H 0.523636 -2.235418 -3.574330 ! O -5.278697 -1.286096 0.570945 O 1.664670 -2.878599 2.093566 ! H -5.578501 -2.001817 -0.011114 H 2.634761 -2.906878 2.032006 ! H -4.494344 -0.879326 0.112135 O 0.148858 -1.108299 3.216148 ! O -6.492384 1.256920 0.187783 H 0.800735 -1.468735 3.840159 ! H -5.629673 1.544062 -0.199551 O 2.474445 -0.130398 1.787251 ! H -6.313336 0.317986 0.407482 H 2.401945 0.662400 1.173329 ! O -1.686892 4.874525 0.318616 H 0.706141 -0.793882 -2.913335 ! H -1.448395 3.937653 0.518668 H -2.259711 -1.279874 -1.263913 ! H -0.932255 5.397875 0.629684 H -0.029433 -0.156996 3.448812 ! O -0.680391 1.492565 3.788480 H 3.325306 -0.540916 1.557918 ! H -0.780698 1.821066 2.859020 C 0.648387 2.256683 -1.524254 ! H -0.063324 2.104800 4.219461 H 0.369720 3.242760 -1.135112 ! H 0.593786 2.289335 -2.616610 ! ! ! ! ! C -1.765181 1.602424 -1.673374 ! ! ! ! ! H -1.738076 1.222460 -2.703176 ! ! ! ! ! H -1.841740 2.690408 -1.729954 ! ! ! ! ! C -3.055107 1.099818 -1.016939 ! ! ! ! ! O -3.124503 -0.134531 -0.663896 ! ! ! ! ! O -3.996697 1.919349 -0.878992 ! ! ! ! ! C 2.092394 1.961759 -1.107641 ! ! ! ! ! O 2.359711 1.983327 0.144987 ! ! ! ! ! O 2.939339 1.758277 -2.019627 ! ! ! ! ! ! ! ! ! !
183 References
Silva,'A.M.N.,'Kong,'X.,'Hider,'R.C.,'2009.'Determination'of'the'pKa'value'of'the' hydroxyl'group'in'the'αDhydroxycarboxylates'citrate,'malate'and'lactate'by'13C' NMR:'implications'for'metal'coordination'in'biological'systems.'BioMetals'22,'771D 778.'
Stumm,'W.,'Morgan,'J.,'1996.'Aquatic'Chemistry'(3rd'edn.'ed.).'Wiley,'New'York.'
184
Appendix B: Crystal description and solution speciation diagrams
185
(A)
(B)
(C)
Figure B.1. Cross-section with the active surface on the top (Left) and overhead view (Right) of (A) (001) surface, (B) (012) surface and (C) (110) surface views of the ideal oxygen truncated hematite surfaces. The directions of the crystallographic axes are also shown for reference. The surfaces vary in the proportions of singly, doubly and triply coordinated oxygen atoms as well as the surface topography.
186 1
0.9
0.8
0.7
0.6
0.5 Fraction 0.4 Pb+2 Pb(OH) (aq) 0.3 2 Pb(OH)3− Pb OH+3 0.2 2 Pb (OH)+2 3 4 Pb (OH)+4 0.1 4 4 PbOH+ 0 2 3 4 5 6 7 8 9 10 11 pH Figure B.2. Lead (0.1 mM) speciation in the a (0.1M NaClO4) background electrolyte calculated by Visual Minteq3.0
187 1
Pb+2 0.9 Pb(OH) (aq) 2 3− 0.8 Pb(OH) Pb OH+3 2 Pb (OH)+2 0.7 3 4 Pb (OH)+4 4 4 0.6 PbOH+ Pb(Citrate)−4 2 0.5 Pb−Citrate−
Fraction PbH(Citrate)−3 0.4 2 PbHCitrate (aq) PbH Citrate+ 0.3 2
0.2
0.1
0 2 3 4 5 6 7 8 9 10 11 12 13 pH Figure B.3. Lead (0.1 mM) Speciation in the presence of 1 mM Citrate in a (0.1M NaClO4) background electrolyte calculated by Visual Minteq 3.0.
188 1
0.9
0.8 Pb+2 Pb(OH) (aq) 0.7 2 Pb(OH)− 3 0.6 Pb OH+3 2 Pb (OH)+2 0.5 3 4 Pb (OH)+4 4 4 Fraction 0.4 PbOH+ Pb−(Phthalate)−2 2 0.3 Pb−Phthalate (aq) PbH−(Phthalate)− 2 0.2 PbH−Phthalate+
0.1
0 1 2 3 4 5 6 7 8 9 10 11 pH Figure B.4. Lead (0.1 mM) Speciation in the presence of 1 mM Phthalate in a (0.1M NaClO4) background electrolyte calculated by Visual Minteq 3.0
189
Appendix C: X-ray Reflectivity
C.1 Setup
Two types of X-ray reflectivity data were collected and analyzed, nonresonant X-ray reflectivity (XR) and resonant anomalous X-ray reflectivity (RAXR). The physical setup for both methods is the same. After equilibrating in the sample solution, the crystal was mounted in a custom-designed flow through cell that was sealed with a Kapton film window shown in figure C1. The cell was flushed with the sample solution and allowed to equilibrate. The drainage port on the cell was left open during this time in order to remove excess sample solution to minimize the absorbance of X-rays by the solution. It is estimated that the thickness of the solution film on the sample is 2 µm (Fenter and
Sturchio, 2004). The sample holder was then placed in the diffractometer.
Alignment of the sample requires several steps. The first step involves aligning the crystal in the vertical direction cutting the incident beam in half and then the theta angle
(rotation around the axis perpendicular to the beam) to make the crystal parallel to the incident X-ray beam. Several iterations were required for proper alignment. As the crystals were all miscut to varying degrees, the crystal had to be aligned by the angle perpendicular to the incident beam, χ, and the rotation about the z axis, perpendicular to the crystal surface, ϕ such that both miscut reflections were visible and centered as shown
190 in figure C1. Once set, ϕ and χ were not changed during the measurement. The detector in this setup is located on the detector arm and the angle is changed at the rate of twice θ.
C.2 Data Collection
Each data point was collected as an image files using either a CCD or a Pilatus detector.
To collect the XR data, the incident beam energy was held constant at 12 keV, a distance away from the binding energy of any lead electrons. The data was collected as a function of L, in reciprocal lattice units, which relates to the electron momentum transfer (q) by:
2! ! = ! !!!!!!!!!!". 1 !
Where d is the crystal lattice spacing. As the lattice spacing changes with the exposed surface and the data was collected in terms of L, the ranges over which q were collected are similar, though not the same for the different surfaces. The range over which q was collected is displayed in table C1. The stability of the system (ie , how much the system changes with X-ray exposure) was measured by either returning to a specific point (on beamline 33-ID, (001) surface only, q = 1.23 and q = 2.65) or collecting every other data point from high q to low q and filling in the missing data points from low q to high q (on beamline 6-ID). The reflectivity from the (001) surface was much less stable than the other two surfaces implying the lead was mobile,(Lee et al., 2011) so the sample was moved approximately every 10 minutes during the collection of a XR scan to illuminate a fresh spot on the crystal. The length of an XR scan depends on the beamline setup (ie. detector, diffractometer speed, flux, etc.) and the reflectivity of the crystal surface with the adsorbates. Generally speaking, XR scan time is between 30 minutes to one hour.
191 For the RAXR measurements, the energy was varied over the lead LIII edge, 13.035 keV
(±300 eV) while holding q constant. This was repeated at multiple values of q from ~1
Å-1 to ~4 Å-1 by which point the edge jump was not apparent. Stability was measured by returning to a single q value and collecting the same spectrum every 15 minutes. If the spectrum changed shape and / or intensity, the sample was flushed and moved to illuminate a new spot. The scan times for each RAXR spectrum varied depending on the signal and the detector used. Some scans required 20 seconds per point while others required only 0.5 seconds.
C.3 XR Modeling
The single crystal nonresonant X-ray reflectivity experiments are a result of the crystal truncation rod (CTR). The CTR arises when the crystal is cut to a specific plane and the scattering is no longer isotropic, but results in a streak in the direction normal to the surface.(Als-Nielsen and McMorrow, 2011) The intensity of the reflection off the single crystal was modeled as:
! 4!!! ! ! ! = !!"!!"# + !!"# + !! !!!!!!!!!!!!!!"!2 !!!"
5 where re is the classical electron radius (2.818x10 Å), AUC is the area of the unit cell of the exposed surface given in table C2, FUC, FCTR, Fint, and FW are the structure factors of the single unit cell of the specific surface being studied, the semi-infinite crystal cleaved to the specific surface, the interfacial region, and the bulk water respectively.(Lee et al.,
2011) The FCTR is given as:
192 1 !!"# = ! !!"# !!!!!!!!!!!!!!!!!!!!!!!!"!3 1 − exp !
The values for d are given in table C2. The structure factors for FUC and Fint are calculated based on:
! −!! ! = ! ! ! ! exp !"! exp ! !!!!!"!4 ! ! ! 2 !
th where fj(q) is the atomic scattering factor of the j atom which was approximated from the nonresonant atomic form factor calculated from the data in the International Table for Crystallography vol. 3 (Wilson, 1992). cj, zj, and uj are the occupancy, location, and root mean square (rms) width of the Gaussian distributions used to describe the individual, jth, atoms. These include both the atoms in the crystal structure and the adsorbed atoms. For the interfacial structure factor, the model was fit using a water equivalent fit where the electron density was calculated assuming that water was the only contributor to the electron density.(Lee et al., 2007) The model used to fit the water was a layered model described by:
! ! ! ! !"# −0.5 !! ! ! = !"#$% !" ! ! ! !!!!!!!!"!5 ! ! 1 − !"# −0.5 !!!"# !"# !"!!
Where fwater is the scattering factor of a water molecule, dw is the layer spacing of water,
-3 ρw is the density of water (0.33 Å ), and σ0 is the rms width of the water interfacial profile.(Fenter and Sturchio, 2004) Equations 2-5 show us that the measured data, R(q), is proportional to the Fourier transform of the electron density with respect to height about the surface. An example of the XR and the fit are given in figure C2.
193 Example input files for the three surfaces are given in tables C3-C5. Each surface required a different model, as the FUC is different for each surface. The (001) surface was fit by defining the unit cell as 6 Fe Fe O layers with occupancies of 1, 1, and 3 respectively. The data was fit by allowing each atom to vary independently. This was not possible with the other two surfaces. On both the (110) and (012) surfaces, neighboring atoms needed to be “bundled” together (e.g., FeO) as the covariance between the individual atoms was too high otherwise.
In the model parameter files (Tables C3 – C5), the first parameter defines what the water distribution will be. This parameter was chosen at the beginning and then was not varied
(as defined by the value in column 3, “1” for fit, “0” for not fit) during the fit. The second parameter is a scaling factor for the whole XR plot. The 3rd parameter is the
Robinson roughness. This parameter attempts to account for surface roughness.(Robinson, 1986) Parameters 5-8 define the water layer as given in eq. 5. As indicated by the first column of the parameter files, several parameters that were not included in the model were not included in the tables. The parameters that define the locations with respect to the surface of each of the atom layers in the crystal are given by parameters 16-27 in the (110), 15-26 in the (012), and 15-32 in the (001) model. As discussed earlier these are different for each surface. All of the layers were allowed to vary independently on the (001) surface. This was not possible on the (110) or (012) as there was a high level of covariance between the neighboring layers, so they were grouped as explained in the parameter name column. The next three parameters after the bulk crystal in each of the parameter files controls the rms width of the top three layers.
194 All of the other bulk crystal layers had a fixed (ie column 3 = 0) rms width. These parameters were generally only allowed to vary at the end of the fitting process and were restricted to only increasing the rms width. Following these parameters, the (012) has the additional half layer parameter to account for the double terminations (parameter x, Table y). It was determined to be ~0.75 from the fit of DI water and allowed to vary only at the end. Additional termination fits were attempted on the other surfaces, but they were not necessary and thus ….. The final 15 parameters in each of the models control the adsorbates. The first 5 (del#_ads) are the position of the Gaussian distribution, the second five (occ#_ads) are the occupancy of the distribution and third five (u_#) are the rms width. The model can use any number of adsorbate peaks from 0 to 5, and individual peaks are removed from the fit by setting the occupancy to 0.
C.4 RAXR Modeling
The RAXR data was collected as a function of energy, E, for each individual scan, and q, over the course of an entire data set (Figure C4). This yields a resonant structure factor,
FR(q0,E), for each scan. The resonant structure factor is split into an energy dependent part and a q dependent part and is given by the following equation (Lee et al., 2011):
! −!!!",! !! !!, ! = ! !!′ ! !+ !!"" ! !!",! exp !"!!",! exp !!!!"!6 ! 2
The energy dependent part is given by !′ ! !+ !!"" ! and it was determined by measuring the X-ray near edge adsorption spectrum of a 0.1 M PbClO4 solution in transmission mode using an ion chamber detector. This was measured only once per 195 beamtime and used for all data collected during that beamtime. The model fitting was done with the q dependent part of the resonant structure factor. The fitting parameters c,Me,J, zMe,J, and uMe,J are the occupancy, height above the hematite surface, and rms width of each lead peak, which was modeled as a Gaussian distribution. The FR(qo,E) is combined with the nonresonant structure factor, FNR(q0), to arrive at the total structure factor. The FNR(q0) was determined from the best fit model of the nonresonant XR data.
A model independent (MI) fitting method(Park and Fenter, 2007) was used to determine the initial model parameters for the model dependent fit (shown in figure C3). The MI fit was accomplished by altering eq. 6 to be:
!! !!, ! = ! !!′ ! !+ !!"" ! A!exp !!! ! !!!!!!!"!7
Where AR is the amplitude and ΦR is the phase of each individual RAXR spectrum. A good fit is achieved when the model dependent phases and amplitudes match the model independent phases and amplitudes as shown in Figure C5.
Both the XR and RAXR data rely on the χ2 value to determine the best fit as given by:
! !! − !!"#$ ! !! !! = ! ! !!!!!!!!!!!!!!!!!!!!"!8 ! − !!
Where Ik and Icalc are the measured and modeled intensities respectively, σk is the uncertainty of the kth data point, and N and Np are the number of data points and parameters in the model fit.(Lee et al., 2011)
196
Table C.1. Range of q over which data was collected for the XR experiments on each surface
-1 -1 Surface qmin (Å ) qmax (Å ) (001) 0.06 5.85 (012) 0.17 5.64 (110) 0.25 5.74
197
Table C.2. Physical parameters of the three surfaces. The d spacing was found to vary slightly between beamtimes and is given here to the maximum precision all the measured values agree to. When the lattice height is off even to the fourth decimal of precision, the Bragg peak location will change and can be seen in data points near the Bragg peak.
2 Surface AUC (Å ) D (Å) (001) 21.95 13.76 (012) 13.66 3.68 (110) 39.96 2.52
198 Table C.3. Input parameter file for (012) model.
No Best Fit Precisio Ma Initial Std Param Name Fit ? n x De 1 2 0 0.0001 St0.1ep 1 v 0 water_model (0:none,1:featureless,2:layered) 2 1 1 0.0001 0.1 1 0 scale 3 0.05 0 0.0001 0.1 0.05 0 beta (robinson roughness) 4 0 0 0.0001 0.1 0 0 omega (single layer roughness) 5 10 0 0.0001 1 10 0 water_thickness 6 9 0 0.0001 0.1 9 0 r0 (water profile height) 7 1.2 0 0.0001 0.1 1.2 0 u0 (water profile width) 8 3.5 0 0.0001 0.1 3.5 0 dwater 12 1 0 0.0001 0.5 0 0 cov_topO 15 1 0 0.0001 0.1 1 0 relative_displacement of FeO in top-3 layer 16 1 0 0.0001 0.1 1 0 relative_displacement of OFe in top-3 layer 17 1 0 0.0001 0.1 1 0 relative_displacement of O in top-3 layer 18 1 0 0.0001 0.1 1 0 relative_displacement of FeO in top-2 layer 19 1 0 0.0001 0.1 1 0 relative_displacement of OFe in top-2 layer 20 1 0 0.0001 0.1 1 0 relative_displacement of O in top-2 layer 21 1 0 0.0001 0.1 1 0 relative_displacement of FeO in top-1 layer 22 1 0 0.0001 0.1 1 0 relative_displacement of OFe in top-1 layer 23 1 1 0.0001 0.1 1 0 relative_displacement of O in top-1 layer 24 1 1 0.0001 0.1 1 0 relative_displacement of FeO in top layer 25 1 1 0.0001 0.1 1 0 relative_displacement of OFeO in top layer 26 1 0 0.0001 0.1 1 0 relative_displacement of only O in top layer 27 1 0 0.0001 0.1 1 0 Proportional vibrational amplitude enhancement for FeO in the top 28 1 0 0.0001 0.1 1 0 Proportionallayer vibrational amplitude enhancement for OFe in the top 29 1 0 0.0001 0.1 1 0 Proportionallayer vibrational amplitude enhancement for O in the top 30 0.75 0 0.0001 0.1 1 0 coveragelayer of the top FeO (i.e., Term2) 31 2 1 0.0001 0.1 2 0 del1_ads: Adsorbate 1-5 position 32 4 1 0.0001 0.1 4 0 del2_ads 33 6 1 0.0001 0.1 6 0 del3_ads 34 8 0 0.0001 0.1 8 0 del4_ads 35 10 0 0.0001 0.1 10 0 del5_ads 36 2 1 0.0001 0.1 2 0 occ1_ads: Adsorbate 1-5 occupation 37 2 1 0.0001 0.1 2 0 occ2_ads 38 2 1 0.0001 0.1 2 0 occ3_ads 39 0 0 0.0001 0.1 0 0 occ4_ads 40 0 0 0.0001 0.1 0 0 occ5_ads 41 1 0 0.0001 0.1 1 0 u_1:: Adsorbate 1-5 rms width 42 1 0 0.0001 0.1 1 0 u_2 43 1 0 0.0001 0.1 1 0 u_3 44 1 0 0.0001 0.1 1 0 u_4 45 1 0 0.0001 0.1 1 0 u_5
199 Table C.4. Input parameter file for (001) model
No Best Fit Fit? Precision Max Initial Std Param Name Step Dev 1 2 0 0.0001 0.1 2 0 water_model (0:none,1:featureless,2:layered) 2 1 1 0.0001 0.1 1 0 scale 3 0.05 1 0.0001 0.1 0.05 0 beta (robinson roughness) 4 0 0 0.0001 0.1 0 0 omega (single layer roughness) 5 10 0 0.0001 1 10 0 water_thickness 6 9 0 0.0001 0.1 9 0 r0 (water profile height) 7 1.3 0 0.0001 0.1 1.3 0 u0 (water profile width) 8 3.5 0 0.0001 0.1 3.5 0 dwater 15 1 0 0.0001 0.1 1 0 relative_displacement of Fe(lower) in the top-5 layer 16 1 0 0.0001 0.1 1 0 relative_displacement of Fe(upper) in the top-5 layer 17 1 0 0.0001 0.1 1 0 relative_displacement of O in the top-5 layer 18 1 0 0.0001 0.1 1 0 relative_displacement of Fe(lower) in the top-4 layer 19 1 0 0.0001 0.1 1 0 relative_displacement of Fe(upper) in the top-4 layer 20 1 0 0.0001 0.1 1 0 relative_displacement of O in the top-4 layer 21 1 0 0.0001 0.1 1 0 relative_displacement of Fe(lower) in the top-3 layer 22 1 0 0.0001 0.1 1 0 relative_displacement of Fe(upper) in the top-3 layer 23 1 0 0.0001 0.1 1 0 relative_displacement of O in the top-3 layer 24 1 0 0.0001 0.1 1 0 relative_displacement of Fe(lower) in the top-2 layer 25 1 0 0.0001 0.1 1 0 relative_displacement of Fe(upper) in the top-2 layer 26 1 0 0.0001 0.1 1 0 relative_displacement of O in the top-2 layer 27 1 0 0.0001 0.1 1 0 relative_displacement of Fe(lower) in the top-1 layer 28 1 0 0.0001 0.1 1 0 relative_displacement of Fe(upper) in the top-1 layer 29 1 0 0.0001 0.1 1 0 relative_displacement of O in the top-1 layer 30 1 1 0.0001 0.1 1 0 relative_displacement of Fe(lower) in the top-0 layer 31 1 1 0.0001 0.1 1 0 relative_displacement of Fe(upper) in the top-0 layer 32 1 1 0.0001 0.1 1 0 relative_displacement of O in the top-0 layer 33 1 0 0.0001 0.1 1 0 uenh_Fe_0L: Proportional vibrational amplitude enhancement for Fe and O in the top Layer 34 1 0 0.0001 0.1 1 0 uenh_Fe_0U 35 1 0 0.0001 0.1 1 0 uenh_O_0 36 2 1 0.0001 0.1 2 0 del1_ads: Adsorbate 1-5 position 37 4 1 0.0001 0.1 4 0 del2_ads 38 6 0 0.0001 0.1 6 0 del3_ads 39 8 0 0.0001 0.1 8 0 del4_ads 40 10 0 0.0001 0.1 10 0 del5_ads 41 2 1 0.0001 0.1 2 0 occ1_ads: Adsorbate 1-5 occupation 42 2 1 0.0001 0.1 2 0 occ2_ads 43 0 0 0.0001 0.1 0 0 occ3_ads 44 0 0 0.0001 0.1 0 0 occ4_ads 45 0 0 0.0001 0.1 0 0 occ5_ads 46 1 1 0.0001 0.1 1 0 u_1:: Adsorbate 1-5 rms width 47 1 1 0.0001 0.1 1 0 u_2 48 1 0 0.0001 0.1 1 0 u_3 49 1 0 0.0001 0.1 1 0 u_4 50 1 0 0.0001 0.1 1 0 u_5
200 Table C.5. Input parameter file for (110) model. No Best Fit? Precision Max Initial Std Param Name Fit Step Dev 1 2 0 0.0001 0.1 2 0 water_model (0:none,1:featureless,2:layered) 2 1 1 0.0001 0.1 1 0 scale 3 0.05 1 0.0001 0.1 0.05 0 beta (robinson roughness) 4 0 0 0.0001 0.1 0 0 omega (single layer roughness) 5 10 0 0.0001 1 10 0 water_thickness 6 9 0 0.0001 0.1 9 0 r0 (water profile height) 7 3.5 0 0.0001 0.1 3.5 0 u0 (water profile width) 8 1.5 0 0.0001 0.1 1.5 0 dwater 9 0.25 0 0.0001 0.1 0.25 0 ubar 16 1 0 0.0001 0.01 1 0 relative_displacement of Fe in the top-3 layer 17 1 0 0.0001 0.01 1 0 relative_displacement of O 43 in the top-3 layer 18 1 0 0.0001 0.01 1 0 relative_displacement of O 21 in the top-3 layer 19 1 0 0.0001 0.01 1 0 relative_displacement of Fe in the top-2 layer 20 1 0 0.0001 0.01 1 0 relative_displacement of O 43 in the top-2 layer 21 1 0 0.0001 0.01 1 0 relative_displacement of O 21 in the top-2 layer 22 1 0 0.0001 0.01 1 0 relative_displacement of Fe in the top-1 layer 23 1 0 0.0001 0.01 1 0 relative_displacement of O 43 in the top-1 layer 24 1 0 0.0001 0.01 1 0 relative_displacement of O 21 in the top-1 layer 25 1 1 0.0001 0.01 1 0 relative_displacement of Fe in the top layer 26 1 1 0.0001 0.01 1 0 relative_displacement of O 43 in the top layer 27 1 1 0.0001 0.01 1 0 relative_displacement of O 21 in the top layer 28 1 0 0.0001 0.01 1 0 uenh_Fe_0: Proportional vibrational amplitude enhancement for Fe and O 29 1 0 0.0001 0.01 1 0 uenh_O43_0in the top layer 30 1 0 0.0001 0.01 1 0 uenh_O21_0 31 2 1 0.0001 0.1 2 0 del1_ads: Adsorbate 1-5 position 32 4 1 0.0001 0.1 4 0 del2_ads 33 6 0 0.0001 0.1 6 0 del3_ads 34 8 0 0.0001 0.1 8 0 del4_ads 35 10 0 0.0001 0.1 10 0 del5_ads 36 3 1 0.0001 0.1 3 0 occ1_ads: Adsorbate 1-5 occupation 37 3 1 0.0001 0.1 3 0 occ2_ads 38 0 0 0.0001 0.1 0 0 occ3_ads 39 0 0 0.0001 0.1 0 0 occ4_ads 40 0 0 0.0001 0.1 0 0 occ5_ads 41 1 0 0.0001 0.1 1 0 u_1:: Adsorbate 1-5 rms width 42 1 0 0.0001 0.1 1 0 u_2 43 1 0 0.0001 0.1 1 0 u_3 44 1 0 0.0001 0.1 1 0 u_4 45 1 0 0.0001 0.1 1 0 u_5
201
A) B)
The Journal of Physical Chemistry C Article
Figure 1. Schematic of the X-ray thin film cell used to collect the data presented in this study. The cell is sealed with a ∼8 μm thick polyimide membrane, and the solution can be exchanged using a syringe. Typical sample dimensions: ∼10 × 10 × 1 mm3.
Figure C.1. A) schematic of the sample cell used for XR and RAXR experiments from National(Bellucci Laboratory. et al., All 2015) measurements B) image were of conducted the (001) in situ hematite sample mounted in the cell on the at room temperature. The beam size at the sample was 150 × 1000diffractometer.μm2 (v × h) with The a flux (012) of ∼5 × and1011 photons/s.(110) crystal The data were 1 cm square rather than 1 cm by 3 cm. were collected with a charge-coupled device (CCD) area detector. From each CCD image, the background was interpolated across the signal region using either a linear or polynomial regression line (chosen by the goodness of the fit to the background shape), and subtracted to isolate the reflectivity signal whose uncertainty was determined by counting statistics.31 In spite of our extensive sample preparation, the quartz surfaces were found to be topographically heteroge- Figure 2. Section view of the quartz(101) structure projected along fl neous. To address this issue, 2D maps of the re ectivity signal the a2 axis. Thin black parallelogram shows a projection of one unit as a function of the beam footprint position were collected at a cell onto the (010) plane. Black dot-dashed lines indicate distinct fixed scattering condition near the first midzone. Areas of the layers and intersect the Si2 atoms. Redrawn after Schlegel et al. 33 sample that showed high and uniform reflectivity, correspond- (2002), with the origin for reference to atom heights chosen as the ing to areas with minimal roughness, were used to collect the average height of the O5 and O6 terminal oxygens (blue dotted line). data presented here. The stabilities of the interfacial systems were monitored by repeated measurements of XR and RAXR at selected q conditions. For the Rb data, the reversibility of Rb fl The data were analyzed with a parametrized electron density adsorption reactions was demonstrated by ushing the sample fi cell with 50 mL ultrapure DIW and observing that no Rb model that consists of the following: (a) the unmodi ed resonant signal was present after flushing (data not shown). subsurface crystal lattice; (b) an interfacial region that includes 2.3.1. X-ray Reflectivity (XR). The XR signal, defined as the the relaxed quartz layers at the surface (the top three quartz ratio of the reflected to incident X-ray flux, was measured at the unit cells for this study, Figure 2), the terminal oxygens (O5 quartz(101) − DIW and the quartz(101)−RbCl solution and O6), and the surface-adsorbed species; and (c) the bulk interfaces as a function of the momentum transfer (q), defined solution above the surface. The origin for the reported atom as q = (2π/d)L where d = 3.3434 Å is the quartz(101) spacing heights (z = 0) is the average height of the unrelaxed O5 and and L is the Bragg index. The reflectivity R(q) can be expressed O6 positions, and all relaxations are reported as the deviations as from the bulk crystallographic positions. The vertical relaxation 22 2 in the quartz surface was modeled by allowing Si1, Si2, Si3, and R()qBqrqAFFFF=| () | (4π eUCUCCTR / ) | +int + w | (1) the terminal O5 and O6 to relax independently while restricting −5 the vertical displacements of the remaining O atoms to be the where re = 2.818 × 10 Å is the classic electron radius, and AUC = 33.8 Å2 is the area of a unit cell along the (101) plane. F is a average displacement of its two coordinated Si atoms. The structure factor [for the substrate bulk unit cell (UC), interfacial occupancy factors of the Si atoms as well as their coordinated fl fi oxygen atoms in the top quartz unit-cell layer (Figure 2) were region (int), or the uid water (w)] de ned as F =Σj oj f j(q) 2 202 allowed to vary to simulate surface defects (e.g., atom exp(iqzj) exp[−(quj) /2], where o is the occupancy, z is the height, u is the vibrational amplitude, and f(q) is the atomic vacancies) which can be present at the natural surfaces used fi scattering factor of each atom, and the summation is over all in the current experiments. The electron density pro le of the fi atoms, j, within a speci c substructure. FCTR = 1/(1 − interfacial region is derived as the sum of the Gaussian exp(−iqd/2)) is the form factor for a semi-infinite crystal distributions for each atomic layer, with parametrized along the surface normal direction (Fenter, 2002). The occupancies, widths, and heights that are defined by the prefactor, |B(q)|2 = (1-β)2/[1+ β2-2βcos(qd)] where β (0 ≤ model fits. Various models for the interfacial water structure β ≤ 1) is the Robinson roughness factor,29,31 accounts for the were explored (e.g., with 1−4 layers modeled by Gaussian q-dependent reduction in the reflectivity signal associated with functions followed by an error function profile extending surface roughness. The XR data were collected at a photon beyond the last layer of structured water). The best fit model energy of 18 keV and for q values between 0.1 and 5.8 Å−1, had two adsorbed water layers. corresponding to a range of Bragg index L from 0.05 to 3.09 The model parameters were optimized through a nonlinear reciprocal lattice units (rlu). The experimental resolution of the least-squares algorithm, and the goodness of fit was evaluated 31 2 fi data (π/qmax) is ∼0.55 Å for both data. through the Pearson’s chi-squared (χ ) value de ned as
4780 DOI: 10.1021/jp510139t J. Phys. Chem. C 2015, 119, 4778−4788 Scan#=35 Image#=55 L=1.46 q=3.642 Å−1 E= 12keV 200 2.72 250 100 300 2.7 200 350 2.68 300 400 2.66 450 400
Y (pixel) 2.64 500 500 550 2.62 600 600 2.6 700 650 2.58 700 800 2 2.5 3 100 200 300 400 500 600 700 800 4 x 10 X (pixel) 4 x 10 1.2
1
0.8 100 200 300 400 500 600 700 Figure C.2. Image taken by the CCD detector of the reflection off the (110) hematite surface. The two reflections are present due to the miscut in the crystal. The overlap of the two rectangles in the area is integrated to get the reflection and the area inside the rectangles, but outside of the overlap is used to determine the background. The dark blue around the edges is the result of the slits placed before the detector. These slits closed more when taking images around the Bragg peak. If the reflection at the Bragg peak hits the detector, it could possibly damage the detector.
203 −4 10 Ideal Term Data Fit −6 Q 10 RAXR Intensity −8 10
−10 10 0 1 2 3 4 5 6 −6 10
−7 10
−8 10
−9 10 0 1 2 3 4 5 6
2
0
−2
Residuals (chi2) −4 0 1 2 3 4 5 6 −1 Q (Å )
Figure C.3. Example of the XR Data collected on the 110 surface. The top plot shows the raw data with empty red circle. The green line is the ideal termination if no adsorbents were present. The black line is the fit to the data. The solid blue circles are the locations where the structure factor was taken for the RAXR fit. The middle plot shows the data normalized to the generic CTR. The bottom plot shows the residuals.
204 Q= 0.25, L= 0.1 Q= 0.31, L= 0.13 Q= 0.37, L= 0.15 Q= 0.5, L= 0.2
10 10 10 10 5 5 5 0 0 0 0 −5 −5 −5 −10 −10 −10 −10
12.6 12.8 13 13.2 13.4 12.6 12.8 13 13.2 13.4 12.6 12.8 13 13.2 13.4 12.6 12.8 13 13.2 13.4
Q= 0.62, L= 0.25 Q= 0.75, L= 0.3 Q= 0.87, L= 0.35 Q= 1.12, L= 0.45
10 5 5 5 5
0 0 0 0
−5 −5 −5 −10 −5 12.6 12.8 13 13.2 13.4 12.6 12.8 13 13.2 13.4 12.6 12.8 13 13.2 13.4 12.6 12.8 13 13.2 13.4
Q= 1.27, L= 0.51 Q= 1.5, L= 0.6 Q= 1.75, L= 0.7 Q= 2, L= 0.8 5 4 4 5 2 2
0 0 0 0
−2 −2 −4 −5 −5 −4 12.6 12.8 13 13.2 13.4 12.6 12.8 13 13.2 13.4 12.6 12.8 13 13.2 13.4 12.6 12.8 13 13.2 13.4
Q= 2.24, L= 0.9 Q= 2.87, L= 1.15 Q= 3.24, L= 1.3 Q= 3.62, L= 1.45 5 10 5 2 5 0 0 0 0 −5 −5 −2 −10 −5 12.6 12.8 13 13.2 13.4 12.6 12.8 13 13.2 13.4 12.6 12.8 13 13.2 13.4 12.6 12.8 13 13.2 13.4
Q= 4.36, L= 1.75 5
0
−5 12.6 12.8 13 13.2 13.4
Figure C.4. Example of a model independent fit of the RAXR data on the (110) hematite surface. The red circles show the data and the blue line is the fit. A baseline is applied to both the data and the fit. In this method, each of the individual spectra are fit and the lead location is determined by combining the fit.
205 Q= 0.25, L= 0.1 Q= 0.31, L= 0.13 Q= 0.37, L= 0.15 Q= 0.5, L= 0.2
10 10 10 10 5 0 0 0 0 −5 −10 −10 −10 −10
12.6 12.8 13 13.2 13.4 12.6 12.8 13 13.2 13.4 12.6 12.8 13 13.2 13.4 12.6 12.8 13 13.2 13.4
Q= 0.62, L= 0.25 Q= 0.75, L= 0.3 Q= 0.87, L= 0.35 Q= 1.12, L= 0.45 5 10 5 5 5
0 0 0 0
−5 −5 −5 −10 −5 12.6 12.8 13 13.2 13.4 12.6 12.8 13 13.2 13.4 12.6 12.8 13 13.2 13.4 12.6 12.8 13 13.2 13.4
Q= 1.27, L= 0.51 Q= 1.5, L= 0.6 Q= 1.75, L= 0.7 Q= 2, L= 0.8 10 4 4 5 2 2 5
0 0 0 0
−2 −2 −5 −4 −4 −5 −10 12.6 12.8 13 13.2 13.4 12.6 12.8 13 13.2 13.4 12.6 12.8 13 13.2 13.4 12.6 12.8 13 13.2 13.4
Q= 2.24, L= 0.9 Q= 2.87, L= 1.15 Q= 3.24, L= 1.3 Q= 3.62, L= 1.45
10 10 2 2 5 0 0 0 0 −5 −10 −2 −2 −10
12.6 12.8 13 13.2 13.4 12.6 12.8 13 13.2 13.4 12.6 12.8 13 13.2 13.4 12.6 12.8 13 13.2 13.4
Q= 4.36, L= 1.75
2
0
−2
12.6 12.8 13 13.2 13.4
Figure C.5. Example of a model dependent fit of the RAXR data on the (110) surface. Red circles are the data and blue line is the fit. This is the same data set as in Figure C3. For the model dependent fit, the fit lines are determined from the model of lead locations. The best fit is then determined from the comparison of the data to those fit lines.
206 1
0.8
0.6
0.4
Partial SF Amplitude 0.2
0 0 0.5 1 1.5 2 2.5 3 3.5 4 4.5 5
2.5
2
1.5
1
0.5
Partial SF Phase/q (A) 0
−0.5 0 0.5 1 1.5 2 2.5 3 3.5 4 4.5 5 momentum transfer,Q
Figure C.6. Comparison of the amplitude (top) and phase (bottom) of each RAXR spectrum (e.g., data in Figures C.3 and C.4). The circles are from the model independent fit (figure C.3) and the line is derived from the model dependent fit (figure C.4).
207 References
Als-Nielsen, J., McMorrow, D., 2011. Elements of modern X-ray physics. John Wiley & Sons.
Bellucci, F., Lee, S.S., Kubicki, J.D., Bandura, A., Zhang, Z., Wesolowski, D.J., Fenter, P., 2015. Rb+ Adsorption at the Quartz(101)–Aqueous Interface: Comparison of Resonant Anomalous X-ray Reflectivity with ab Initio Calculations. The Journal of Physical Chemistry C 119, 4778-4788.
Fenter, P., Sturchio, N.C., 2004. Mineral–water interfacial structures revealed by synchrotron X-ray scattering. Progress in Surface Science 77, 171-258.
Lee, S.S., Nagy, K., Park, C., Fenter, P., 2011. Heavy Metal Sorption at the Muscovite (001) -- Fulvic Acid Interface. Environmental Science & Technology.
Lee, S.S., Nagy, K.L., Fenter, P., 2007. Distribution of barium and fulvic acid at the mica-solution interface using in-situ X-ray reflectivity. Geochimica et Cosmochimica Acta 71, 5763-5781.
Park, C., Fenter, P.A., 2007. Phasing of resonant anomalous X-ray reflectivity spectra and direct Fourier synthesis of element-specific partial structures at buried interfaces. Journal of Applied Crystallography 40, 290-301.
Robinson, I., 1986. Crystal truncation rods and surface roughness. Physical Review B 33, 3830.
Wilson, A.J.C., 1992. International tables for crystallography, v. C: Mathematical, physical, and chemical tables. Dordrecht: Kluwer Academic Publishers, Dordrecht.
208
Appendix D: Comparison of (001), (012), and (110) RAXR
As discussed in Chapters 3 and 4, the low amount of lead adsorbed on the hematite surface lead to small edge jumps in the RAXR spectra. Below is a comparison of the fits to the model dependent fits to the RAXR data. In these images, the fit, shown with a blue line is determined by the model and the quality of the fit is determined by how well it matches the data, shown with red circles. The RAXR spectra on the (001) surface have larger errors associated with relative to the size of the edge jump leading to some difficulties in fitting the data as discussed in Chapter 3.
209 Q= 0.41, L= 0.9 Q= 0.55, L= 1.2 Q= 0.69, L= 1.5 Q= 0.82, L= 1.8 1.5 1 1 1 1
0.5 0.5 0.5 0.5
0 0 0 0
−0.5 −0.5 −0.5 −0.5
−1 −1 −1 −1 −1.5 12.6 12.8 13 13.2 13.4 12.6 12.8 13 13.2 13.4 12.6 12.8 13 13.2 13.4 12.6 12.8 13 13.2 13.4
Q= 0.96, L= 2.1 Q= 1.1, L= 2.4 Q= 1.23, L= 2.7 Q= 1.46, L= 3.2 0.4 0.8 0.6 0.4 0.3 0.6 0.4 0.3 0.4 0.2 0.2 0.2 0.2 0.1 0.1 0 0 0 0
−0.2 −0.1 −0.1 −0.2 −0.4 −0.2 −0.2 −0.6 −0.4 −0.3 −0.3 −0.8 −0.4 −0.6 −0.4 12.6 12.8 13 13.2 13.4 12.6 12.8 13 13.2 13.4 12.6 12.8 13 13.2 13.4 12.6 12.8 13 13.2 13.4
Q= 1.74, L= 3.8 Q= 2.1, L= 4.6 Q= 2.47, L= 5.4
0.3 0.2 0.2 0.15 0.2 0.1 0.1 0.1 0.05 0 0 0 −0.05 −0.1 −0.1 −0.1 −0.2 −0.15 −0.3 −0.2 −0.2
12.6 12.8 13 13.2 13.4 12.6 12.8 13 13.2 13.4 12.6 12.8 13 13.2 13.4 Figure D.1. Model fit (blue line) to the RAXR data (red circles) for lead only at pH 6 on the (001) surface.
210 Q= 0.26, L= 0.15 Q= 0.34, L= 0.2 Q= 0.43, L= 0.25 Q= 0.6, L= 0.35 10 4 10 5 5 5 2
0 0 0 0
−5 −5 −2 −5 −10 −10 −4 12.6 12.8 13 13.2 13.4 12.6 12.8 13 13.2 13.4 12.6 12.8 13 13.2 13.4 12.6 12.8 13 13.2 13.4
Q= 0.77, L= 0.45 Q= 0.77, L= 0.45 Q= 0.94, L= 0.55 Q= 1.11, L= 0.65
3 4 3 3 2 2 2 2 1 1 1 0 0 0 0 −1 −1 −1 −2 −2 −2 −2 −3 −3 −3 −4 12.6 12.8 13 13.2 13.4 12.6 12.8 13 13.2 13.4 12.6 12.8 13 13.2 13.4 12.6 12.8 13 13.2 13.4
Q= 1.28, L= 0.75 Q= 1.45, L= 0.85 Q= 1.96, L= 1.15 Q= 2.13, L= 1.25 5 5 5 5
0 0 0 0
−5 −5 −5 −5 12.6 12.8 13 13.2 13.4 12.6 12.8 13 13.2 13.4 12.6 12.8 13 13.2 13.4 12.6 12.8 13 13.2 13.4
Q= 2.31, L= 1.35 Q= 2.48, L= 1.45 Q= 2.82, L= 1.65
5 4 3 2 2 1
0 0 0
−1 −2 −2
−5 −4 −3 12.6 12.8 13 13.2 13.4 12.6 12.8 13 13.2 13.4 12.6 12.8 13 13.2 13.4 Figure D.2. Model fit (blue line) to the RAXR data (red circles) for lead only at pH 6 on the (012) surface.
211 Q= 0.31, L= 0.13 Q= 0.37, L= 0.15 Q= 0.5, L= 0.2 Q= 0.62, L= 0.25 6 6 5 4 5 4 2 2 0 0 0 0
−2 −2 −5 −5 −4 −4 −6 −6 12.6 12.8 13 13.2 13.4 12.6 12.8 13 13.2 13.4 12.6 12.8 13 13.2 13.4 12.6 12.8 13 13.2 13.4
Q= 0.87, L= 0.35 Q= 1.12, L= 0.45 Q= 1.5, L= 0.6 Q= 1.75, L= 0.7 6 3 4 2 4 2 2 2 1 1 0 0 0 0
−2 −1 −1 −2 −4 −2 −2 −6 −3 −4 12.6 12.8 13 13.2 13.4 12.6 12.8 13 13.2 13.4 12.6 12.8 13 13.2 13.4 12.6 12.8 13 13.2 13.4
Q= 2, L= 0.8 Q= 2.24, L= 0.9 Q= 2.87, L= 1.15 Q= 3.24, L= 1.3 15 3 10 5 5 2 5 1 0 0 0 0
−5 −1 −5 −5 −2 −10 −3 −15 12.6 12.8 13 13.2 13.4 12.6 12.8 13 13.2 13.4 12.6 12.8 13 13.2 13.4 12.6 12.8 13 13.2 13.4
Q= 3.62, L= 1.45 Q= 4.36, L= 1.75 6 4 4 2 2 0 0 −2 −2 −4 −4 −6 12.6 12.8 13 13.2 13.4 12.6 12.8 13 13.2 13.4 Figure D.3. Model fit (blue line) to the RAXR data (red circles) for lead only at pH 6 on the (110) surface.
212
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