Geo Chemistry
Geo Chemistry
Geo Chemistry
Georges Calas
Mineral-Aqueous
Solution Interfaces
and Their Impact
on the Environment
Each issue of Geochemical Perspectives presents a single article with an in-depth view
on the past, present and future of a field of
geochemistry, seen through the eyes of highly
respected members of our community. The
articles combine research and history of the
fields development and the scientists opinions
about future directions. We welcome personal
glimpses into the authors scientific life, how
ideas were generated and pitfalls along the way.
Perspectives articles are intended to appeal to
the entire geochemical community, not only to
experts. They are not reviews or monographs;
they go beyond the current state of the art,
providing opinions about future directions and
impact in the field.
Copyright 2012 European Association of Geochemistry,
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ISSN 2223-7755 (print)
ISSN 2224-2759 (online)
DOI 10.7185/geochempersp.1.4
Principal Editor in charge of this issue
Liane G. Benning
Reviewers
Editorial Board
Liane G. Benning
University of Leeds,
United Kingdom
Tim Elliott
University of Bristol,
United Kingdom
Eric H. Oelkers
University of Copenhagen,
Denmark
Marie-Aude Hulshoff
Graphical Advisor
Juan Diego
RodriguezBlanco
University of Leeds, UK
President
Vice-President
Past-President
Treasurer
Secretary
Goldschmidt Officer
Goldschmidt Officer
www.eag.eu.com
CONTENTS
Preface . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . V
Acknowledgments . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . VI
Mineral-Aqueous Solution Interfaces
and Their Impact on the Environment . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 483
Abstract . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 483
1. Introduction . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 485
1.1 A Brief History of Early Geochemical Time . . . . . . . . . . . . . . . . . . 485
1.2 A New Era of Inorganic Geochemistry . . . . . . . . . . . . . . . . . . . . . . . 487
2. Our Personal and Scientific Histories . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 490
2.1 The Life and Times of Gordon Brown . . . . . . . . . . . . . . . . . . . . . . . . 490
2.2. The Life and Times of Georges Calas . . . . . . . . . . . . . . . . . . . . . . . . . 499
3. Historical Perspectives on Crystal Chemistry
and Mineral-Surface Chemistry The Contributions
of Goldschmidt, Pauling, and Gibbs . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 509
3.1. Goldschmidts Atomistic Views of Geochemistry . . . . . . . . . . . . . . 509
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585
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633
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687
688
689
689
690
690
692
692
References . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 695
List of Acronyms . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 733
Index . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 737
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PREFACE
Photos
of (a) Victor Goldschmidt (from Mason, 1992, with permission from the Geochemical
Society). (b) Konrad Krauskopf (source: National Academy of Sciences). (c) Irving Langmuir
(source: National Museum of American History). (d) Linus Pauling (source: Official Nobel Prize
Website). (e) Werner Stumm (source: World Cultural Council).
GEOCHEMICAL PERSPECTIVES | G O R D O N E . B R O W N J R . G E O R G E S C A L A S
VII
Although we also discuss some of the work of others in this area, this
Perspective is not intended to be an exhaustive overview of the field, and we
apologise in advance for omission of the work of many who have contributed to
its development, particularly in the area of microbe-mineral surface interactions,
which is an active and growing research area that should be the subject of a
separate Perspective. In keeping with the objectives of the founders of Geochemical
Perspectives, we offer personal perspectives and discuss how ideas were generated and developed through time, concluding with our thoughts about the future
directions and impact of this field.
Why is it important to understand the molecular-scale factors controlling
geochemical processes? We view this as a key question to address here because of
the disturbing trend in some university Earth science departments to eliminate
undergraduate courses in crystallography, mineralogy, or inorganic geochemistry that focus on molecular-scale concepts and processes. Empirical macroscopic approaches do not lead to a fundamental understanding of the chemical
processes that control Earths near-surface environment or to robust predictive
models of these processes. It is our hope that this Perspective will demonstrate
how molecular-scale geochemical processes occurring at mineral-water interfaces
have a major impact on global geochemical processes.
Gordon E. Brown, Jr.1,2 and Georges Calas3
ACKNOWLEDGMENTS
We wish to thank the US and French research agencies that have generously
supported the research highlighted in this Perspective over the past 30 years,
including the US National Science Foundation (Earth Sciences, Chemistry, and
Biology Divisions) and Department of Energy (Offices of Basic Energy Sciences
and Biological and Environmental Research), the French Centre National de la
Recherche Scientifique (Universe Sciences and Physics Divisions), the Ministry
for Higher Education and Research, the Agence Nationale de la Recherche, the
Commissariat lEnergie Atomique (Direction of the Nuclear Energy), and the
1. Surface & Aqueous Geochemistry Group, Department of Geological and Environmental Sciences,
Stanford University, Stanford, CA, 94305-2115, USA
2. Department of Photon Science and Stanford Synchrotron Radiation Lightsource, SLAC National
Accelerator Laboratory, 2575 Sand Hill Road, MS 69, Menlo Park, CA, 94025, USA
3. Institut de Minralogie et de Physique des Milieux Condenss (IMPMC), UMR 7590, CNRS,
Universit Pierre et Marie Curie, case 115, 4 Place Jussieu, 75252 Paris, Cedex 05, France
VIII
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MINERAL-AQUEOUS
SOLUTION INTERFACES
AND THEIR IMPACT
ON THE ENVIRONMENT
ABSTRACT
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1.
INTRODUCTION
Copyright: NASA.
The interfaces between minerals and aqueous solutions (or atmospheric gases)
play a major role in the geochemistry of Earths surface environment. The importance of such interfaces on a global scale is obvious when Earth is viewed from
space (Fig. 1.1). At the molecular scale, reactions between mineral surfaces and
aqueous solutions, atmospheric gases, and biological organisms are responsible
for chemical weathering and the production of soils. Mineral-surface reactions
also help control the sequestration, release, transport, and transformation of environmental contaminants and pollutants, both natural and anthropogenic, as well
as plant nutrients in the biosphere. In addition, reactions of mineral surfaces with
acid rain are responsible for the degradation of natural and synthetic building
materials. Moreover, mineral carbonation that has sequestered huge quantities
of CO2 naturally over geologic time is a mineral-surface reaction. Werner Stumm
(Stumm et al., 1987) summed up the importance of the mineral-water interface: Almost all the problems associated
with understanding the processes that
control the composition of our environment concern interfaces, above all, the
interfaces of water with naturally occurring solids. The importance of solid
surfaces to chemistry is well summarised in a quotation from Irving Langmuir (Langmuir, 1916): From the point
of view of the chemist, the structure of
the surface must be of utmost importance,
for the chemical reactions in which solids
take part are practically always surface
reactions. Finally, the inherent difficulty encountered in studies of solid
surfaces is captured in a famous Figure 1.1 1972 view of Earth from Apollo
17 emphasising the interfaces
1927 quote from Wolfgang Pauli: God
between continents and oceans
made the bulk, surfaces were invented
(solid-liquid), continents and
by the devil. This may be true, but
atmosphere (solid-gas), and
what great fun mineral surfaces have
oceans and atmosphere (liquidproven to be.
gas).
1
.1 A Brief History of Early Geochemical Time
Inorganic geochemistry has undergone enormous changes since the times of
Victor Moritz Goldschmidt (1888-1947), who spent much of his scientific career
studying the distribution and amounts of chemical elements in minerals, ores,
rocks, soils, waters, the atmosphere, and their natural cycles, on the basis of the
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properties of their atoms and ions (Goldschmidt, 1937). This field, and the closely
related field of mineralogy, have also undergone revolutionary changes during
both of our professional lifetimes due mainly to major advances in analytical
and characterisation methods, theory, and computational power (Brown et al.,
2006a,b). Change is an inevitable and desirable part of life and science, but it
often occurs so rapidly in science that the pioneers in a particular field quickly
become distant memories, or their seminal ideas are not properly introduced to
younger generations of scientists. For example, we have found that some of our
younger colleagues in geochemistry and mineralogy do not recognise the name
Linus Pauling (1901-1994) awarded two unshared Nobel Prizes, one in Chemistry
for elucidating the nature of chemical bonds and the structure of molecules, and
one in Peace for advocating nuclear disarmament and by inference they do
not know of Paulings major contributions to structural chemistry. We have also
found that some of our younger colleagues are unaware of the pioneering studies
in inorganic geochemistry by V.M. Goldschmidt even though they recognise his
name because of the annual international geochemistry conference in his honour.
Similar statements can also be made about Irving Langmuir (1881-1957), who
was awarded the Nobel Prize in Chemistry in 1932 for his theory dealing with
the chemical forces at the boundaries between different substances, or about
Konrad Krauskopf (1910-2003; 1982 Goldschmidt Medalist), one of the pioneers
of inorganic geochemistry in the US, or about Werner Stumm (1924-1999; 1998
Goldschmidt Medalist), one of the pioneers of mineral-water interface chemistry.
This Perspectives issue is, in part, our attempt to change this landscape and show
how the legacy of these pioneers has shaped our field.
Goldschmidts Views on Geochemistry Goldschmidt was not shy in his thinking
about the importance of geochemistry, as illustrated by a quotation from Geochemistry (Goldschmidt, 1958, edited by Alex Muir following Goldschmidts death): From
a human point of view, geochemistry is of the greatest practical importance, especially in
its applications to mining, metallurgy, chemical industry, agriculture and, of course, the
study of terrestrial materials, particularly the accessible outermost parts of our planet
The results of mineralogy, petrology, and geology form the foundation of geochemistry.
Modern atomic chemistry and atomic physics, as well as physical chemistry and chemical physics, give in many cases an essential basis for the understanding of geochemical
problems. Geochemistry, however, is not a debtor only in its relations to theoretical chemistry and physics, since modern inorganic crystal chemistry originated from the study of
geochemical problems, e.g., the practical use of X-ray spectra for chemical analysis and the
development of quantitative optical spectroscopy with the carbon arc. Goldschmidt was
also prescient in his thinking about the linkages among geochemistry, biology, and
the environment: Very close relationships exist between modern geochemistry and pure
and applied biology. The circulation and distribution of many chemical elements in nature
(are) closely related to biochemical processes in which both plants and animals are involved.
Some of the dominant geochemical factors of our time result from the activities of modern
man agriculture, mining, and industry. We resonate particularly with Alex Muirs
and Goldschmidts thoughts on the important connections between observation and
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theory: In the general evolution of geochemistry during the last quarter of a century, the
most remarkable trend is the tendency not only to accumulate analytical facts, but to find a
theoretical explanation of these facts. (Goldschmidt, 1958).
Our own Perspective reflects this idea.
One of us (GB) was fortunate to have known Linus Pauling, Konrad Krauskopf,
and Werner Stumm during his career at Stanford University, and both of our
professional careers have benefitted from their pioneering scientific contributions.
In this Perspective, we discuss some personal interactions with these pioneers
as well as some of their contributions that have impacted our understanding
of mineral-water interface chemistry. We also give an overview of some of our
studies in inorganic geochemistry and mineralogy that are directly related to
the contributions of Pauling, Krauskopf, and Stumm, as well as to those of Goldschmidt and Langmuir. We should not forget that many of the basic principles of
geochemistry, chemical bonding and molecular structure, and surface chemistry
were initially developed by Goldschmidt, Pauling, and Langmuir, respectively. As
part of our discussion, we revisit some of the classical atomistic concepts of crystal
chemistry developed by Pauling (bond valence and Paulings Rules; Pauling, 1927,
1929) and Goldschmidt (ionic radii and the radius ratio principle; Goldschmidt,
1937), and show how they can be extended to mineral-aqueous solution interface
chemistry. We also expand on some of Langmuirs seminal ideas about adsorption
reactions at solid-gas and solid-liquid interfaces that helped define the field of
surface chemistry and provided the framework for our field. Krauskopfs classic
experimental work on the role of adsorption reactions on mineral surfaces and
natural organic matter in controlling trace element concentrations in seawater
(Krauskopf, 1956) is also discussed in light of more recent molecular-level spectroscopic studies (e.g., Brown and Parks, 2001). In addition, we revisit and extend
some of the ideas of Stumm about surface complexation reactions at mineral-water
interfaces (Stumm, 1995) and about chemical weathering (Stumm et al., 1983),
particularly in light of modern spectroscopic and theoretical studies of such reactions. It is our hope that this Perspective will introduce our younger colleagues
in geochemistry and mineralogy, and reintroduce some of our older colleagues,
to these pioneers of our science and will show that their scientific legacy is as
relevant today as it has been over the past 75 years.
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the geometric and electronic structures of adsorbates (e.g., Ford et al., 2001; Brown
and Sturchio, 2002) and the hydrated mineral surfaces to which they attach (e.g.,
Trainor et al., 2004; Yamamoto et al., 2010). Both of us have spent much of our
careers exploiting synchrotron X-ray sources in addressing molecular-level problems in inorganic geochemistry, mineralogy, and most recently environmental
geochemistry and mineralogy (Brown and Calas, 2011; Calas and Brown, 2011).
We have had the benefit of observing first-hand the development of these light
sources and the various synchrotron radiation-based methods they have made
possible in our fields since the first synchrotron radiation facility that allowed
user measurements the Stanford Synchrotron Radiation Project started in
1974 (http://www.slac.stanford.edu/history/ssrp.shtml), followed two years later,
by the first French facility, the Laboratoire pour lUtilisation du Rayonnement
Electromagntique (LURE); http://www.lactualitechimique.org/larevue_article.
php?cle=2607). As both of us have been involved at these centres since these
pioneering times, we will discuss some of the interface-related science these light
sources have made possible as well as some of the interface-sensitive methods
and complementary computational chemistry approaches used to understand
mineral-water interface processes in complex natural environments.
Both of us entered the Earth sciences because of our love of minerals,
our curiosity about how they form, and what they tell us about geological and
geochemical processes. An important part of this Perspective will therefore focus
on environmental mineralogy and geochemistry, which is an area that has
attracted our attention for the past 25 years and which is dominated by chemical
reactions that occur at mineral-water interfaces.
Minerals come in all sizes and shapes, depending on their growth conditions, and range from single crystals many metres in length and girth to nanomaterials with at least one dimension less than 100 nm. A good comparison is
seen in Figure 1.2, which shows giant gypsum crystals found in a lead-zinc-silver
mine in Naica, Chihuahua, Mexico (Garcia-Ruiz et al., 2007), and gypsum crystals
grown in the lab by self-assembly from tiny calcium sulphate precursor nanocrystals (Van Driessche et al., 2012). Because of growing recognition that natural,
incidental, and manufactured nanoparticles play a major role in environmental
processes (Banfield and Navrotsky, 2001; Hochella et al., 2008) and growing
concern that this role is not always positive (Moore, 2006; Choi and Hu, 2008),
we also discuss nanominerals and mineral and mineraloid nanoparticles in the
context of interfacial geochemical processes at Earths surface.
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Figure 1.2
(a) Giant gypsum (CaSO 42H2O) crystals in the Cave of the Crystals in Naica,
Mexico discovered in 2000. Some of these crystals reach lenghts of 11m
(humans for scale), can weigh over 55tons, and are estimated to have taken
up to 0.5 million years to grow. (b) Transmission electron microscope images
of bassanite (CaSO 40.5H2O) nanocrystals (bottom inset) and self-assembled
nanoparticle aggregates (main image) during the formation of gypsum. The
difference in diameter between the giant crystals from Naica and the nanocrystals is about nine orders of magnitude (1-3 m vs. 20nm).
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2.
2
.1 The Life and Times of Gordon Brown
I was born in 1943 in San Diego, California. However, I grew up from age two
in the Deep South of the USA (Mississippi) and experienced first-hand the civil
rights movement of the 1950s and 1960s. As one of my high school classmates
Jim Barksdale (co-founder of Netscape) said in a speech he gave at our 40th high
school reunion in 2001, Things have changed since we grew up in Mississippi and
attended Murrah High School (in Jackson, MS). Rosa Parks sat down (on a bus in
Montgomery, Alabama in 1955), and Sputnick (a Soviet Union spacecraft launched
in 1957) went up. Both of these events profoundly changed our world. Those
were interesting times when the culture in the Deep South was shaken to its
foundations and was changed for the better by these events and others. Three
other events in those early years also helped shape my life: having a 3rd grade
teacher Janet Heredeen who had a simple but magical mineral collection in her
classroom, which began my love of minerals; having a 7th grade science teacher
Nancy Lay who taught me some of the wonders of natural science; and most
important, meeting my future wife of 45 years Nancy Tweedy in that 7th grade
science class (we were married in 1965, and Nancy passed away in 2010).
I attended Millsaps College a liberal arts college in Jackson, MS and
graduated in 1965 with B.S. degrees in Chemistry and Geology. During college,
I was influenced by the small book by William S. Fyfe entitled Geochemistry of
Solids (Fyfe, 1964), which showed how the principles of chemical bonding could
provide useful qualitative insights about the structural chemistry of minerals,
and the book by Linus Pauling entitled The Nature of the Chemical Bond and the
Structure of Molecules and Crystals: An Introduction to Modern Structural Chemistry
(Pauling, 1960), which introduced me to the world of chemical bonding and
structure in molecules and crystals. In part because of these two books and the
encouragement of my advisors in Chemistry (Prof. Roy Berry) and Geology (Prof.
Richard Priddy), I decided to combine these two subjects in my graduate work
and entered the Department of Geochemistry and Mineralogy at Pennsylvania
State University in 1965 to pursue the Ph.D. degree. During my first quarter at
Penn State, I met an Assistant Professor named Jerry Gibbs (Fig. 2.1), who introduced me to X-ray crystallography. Jerry was the best teacher I ever had, and
he became a role model for me professionally and a dear friend. At Penn State
I was also greatly influenced by Prof. Will White, who first sparked my interest
in spectroscopy, and by the late Prof. Rustum Roy, who introduced me to crystal
chemistry and materials science, which was developing as a separate discipline
at that time. I accompanied Jerry Gibbs to Virginia Tech, where he moved in
1966, and completed my M.S. and Ph.D. degrees under his guidance in 1968 and
1970, respectively. During my four years at Virginia Tech, I worked on chemical
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bonding in minerals and the structural chemistry of osumulite and the olivines.
I was also fortunate to meet Charlie Prewitt, then a crystallographer at DuPont
Central Research Labs in Wilmington, Delaware, and learned of his project
with Bob Shannon on updating and quantifying the radii of ions in oxide- and
fluoride-based solids as a function of their coordination number, oxidation state,
and electron spin state. Charlie moved to SUNY Stony Brook in 1969 to establish
a new programme in petrologic crystal chemistry with Jim Papike and invited me
to join them as a postdoctoral student.
During the 16 months I spent at Stony Brook, NY, where my son Michael
was born, I worked on returned Lunar samples, particularly the structural chemistry of Lunar pyroxenes, and how
they could be used to interpret the
thermal history of Lunar basalts,
the development and application
of high-temperature crystal structure methods to minerals, and Al/Si
order-disorder in the feldspars using
neutron diffraction methods.
After a two-year stay (19711973) as an Assistant Professor of
Geological & Geophysical Sciences
at Princeton University, where my
daughter Tracey was born and where
I worked with my first graduate
student George Harlow, I was lured to
Stanford University in 1973 by Dick Figure 2.1 Jerry Gibbs (left) and Gordon
B r o w n (r i g h t ) a t t h e 2 0 0 7
Jahns, Konnie Krauskopf, and Bill Luth,
Geological Society of America
where I have been since. Here, I was
(GSA) Meeting in Denver, CO
introduced to gem-bearing granitic
following acceptance of the
Roebling Medal by Gordon.
pegmatites by Dick Jahns, continued
work on the structure of minerals at
high temperature, and began work on the structural chemistry of silicate glasses
and melts at high temperature with my students Mark Taylor and Mike Hochella
and on feldspars with my first postdoc Phil Fenn. During this same period, my
student Bernard De Jong and I carried out some early semi-empirical quantum
chemical calculations on H6IV T2O7 molecules (T = Si, Al), primitive by todays
standards, X-ray emission spectroscopy studies of glasses on the effects of protons,
hydroxyl groups, CO2, and alkali cations on their energetics and structures in
aluminosilicate glasses and minerals. The results of these studies gave us some
mechanistic insights about the disruptive effects of network modifiers in silicate
glasses and melts. Also at Stanford in 1977, I began using the Stanford Synchrotron Radiation Laboratory (now the Stanford Synchrotron Radiation Lightsource
SSRL) at the Stanford Linear Accelerator Center (SLAC; now officially the SLAC
National Accelerator Laboratory supported by the US Department of Energy
Office of Basic Energy Sciences; DOE-BES) to carry out the first X-ray absorption
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Figure 2.3
Photograph of some of the participants at the 1979 Gordon Research Conference on Inorganic Geochemistry. Front row (all from left): Gordon Brown,
Wes Hildreth, Tony Lasaga, Ian Carmichael, Dean Presnall, Jan Bottinga, Alex
Navrotsky. Second row: Mike Hochella (with hair), Gail Mahood, Enrique
Merino, N. Gray, Tony Philpotts, Georges Calas (looking down on a fly on
Alex Navrotskys head). Third row: Bernie Wood, Bill McKenzie, Charles
Langmuir, John Weare, W. Park, George Flowers. Fourth row: Pascal Richet,
Hans Engi, Richard Sack, J. Wenzel, Tony Morse. Fifth row: E. Takahashi, Bill
Nash, Jim Thompson (1985 Goldschmidt Medalist), S. Hasse, Murli Manghnani,
P.Danckwerth, A. Rite.
When Bill retired from Stanford in 1992, he asked me to take responsibility for one
of his long-time research associates a surface physicist named Tom Kendelewicz,
who had been a mainstay in Bills research group for almost 15 years. Tom was
an expert in synchrotron-based XPS, and he and I wrote a successful proposal to
the US-National Science Foundation (NSF) Earth Sciences Division in 1992 that
resulted in some of the first synchrotron XPS studies of the reaction of water with
mineral surfaces. I recruited a very bright graduate student Ping Liu to work
with us on this project, which provided new insights about this most fundamental
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W hen Alexis
joined my group, she
was aware of our
sy nchrot ron-based
work on the interaction of various cations
and oxoanions at
mineral-water interfaces and wanted to
add the complexity
of microorganisms
to t hese a l ready
complex s y stem s.
I knew little about
microbiology at that
Figure 2.4
(from left to right) Tom Trainor, Peter Eng, Sara time, so I suggested
Trainor, Alexis Templeton, and Mark Rivers at a
that we discuss this
koala preserve south of Melbourne, Australia,
playing hooky one afternoon from the 2006 idea w it h A lf red
Goldschmidt Conference in Melbourne, Australia. Spormann, a recently
a r r ived A ssi sta nt
Professor of Environmental Microbiology at Stanford and an expert on microbial biofilms. Intellectual sparks flew at that first meeting, and we decided to prepare a funding
proposal to initiate this work, which was submitted to the DOE-BES Geosciences
programme. The then-manager of this programme sent the proposal back within
the week without external review indicating that DOE-BES Geosciences didnt
fund that type of work. But we did not give up!
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Ona-Nguema were also members of this team, as were Lars Pettersson (Stockholm
University), Andrea Foster and James Rytuba (U.S. Geological Survey), and Jennifer
Wilcox (Stanford) (Fig. 2.5).
Figure 2.5
Figure 2.6
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I got two years of training for the competition to enter the Grandes Ecoles, a
system that provides French studies of a specific flavour. At the same time, I was
also spending time in the mineralogy collection and laboratory at the Musum
National dHistoire Naturelle de Paris, which was open to undergrads. I was
taking lectures and discussing the history of mineralogy with Jean Orcel, the chair
of mineralogy and successor of Alfred Lacroix in this institution. This experience
helped strengthen my motivation for mineralogy, as my actual studies were
mostly mathematics and physics with some biology and geology. Two years later,
I passed the competition for various higher education institutions during the May
1968 riots. These frequently lasted all night in the streets of Paris, yet in the early
mornings we took exams for several hours in historical buildings that preserved a
miraculous calm during these unpredictable events. I finally chose to enter Ecole
Normale Suprieure (ENS) de Saint-Cloud, a school which later relocated in
Lyon to become the ENS de Lyon. The ENS system is rather unique in the world.
It provides undergrad students a generous stipend during four years, provided
they agree to serve the French public system during 10 years. I was following the
Geology classes at the University of Paris, in its former location in the Sorbonne.
Meanwhile, I was also receiving complementary training at ENS Saint-Cloud in
some fields that were not taught at University, such as computer science, which
was still a minor field. As part of this training, we had geology and ecology field
trips in all parts of France, and in addition I was working during summers as an
adjunct collaborator with the French Geological Survey, which provided me an
unique opportunity to experience the complexity of natural systems.
In 1969, the LMCP moved to the new campus of Jussieu, where it is still
located, and I began several internships in this lab. LMCP was chaired at this time
by Jean Wyart, who encouraged structural studies of new minerals. I was not really
interested in structure determination and decided to move to mineral physics. I
was attracted by books showing other aspects of mineralogy and linking mineralogy and geochemistry. In particular, Geochemistry of Solids (Fyfe, 1964) and
Mineralogical Applications of Crystal Field Theory (Burns, 1970) were enlightening
at this time. I had a fascinating professor in crystal physics, Hubert Curien (19242005), who explained clearly and simply the most complex concepts on defects
in crystals and physical properties of crystalline matter. Most students were so
fascinated by his lectures that they often forgot to take notes, which frequently
caused bad scores on the final exam. Curien, a life fellow of MSA, occupied the
most important managerial and political positions in the French scientific system,
including serving four times as Minister of Research and Technology. I decided to
work on the crystal chemistry of trace elements and radiation-induced defects in
natural fluorites. I was supervised by Roger Maury, an Assistant Professor working
with Curien. Maury was above all a talented pedagogue and introduced me to
mineral spectroscopy and crystal physics. At this time, the laboratory was lucky to
have a branch at one of the best schools in Paris, the Ecole Suprieure de Physique
et de Chimie Industrielle (ESPCI) of Paris, which gained much attention some
years later in having two Nobel Laureates, Pierre-Gilles de Gennes and Georges
Charpak. I was free to use cutting-edge research tools at ESPCI (e.g., an Electron
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within glass structure, and by Guillaume Ferlat and Gerald Lelong (both Assistant
Professors at UPMC in Paris) who carried out mesoscale numerical modelling of
glass structure and high-pressure modifications of glasses, respectively. Other
directions of research concern the structure-property relationships of glasses
and of silicate melts, after pioneering experiments at SSRL with Gordon Brown,
Franois Farges, and Laurence Galoisy, and more recent activities at IMPMC, on
the structural evolution of glasses at high-pressure and the nucleation of glasses.
Determining the surroundings of transition elements in silicate glasses and melts
was also helpful for rationalising mineral-melt partition coefficients. Together
with Gordon, we also share an interest in the location of impurities in minerals,
though with limited activity due to the lack of funding. After the early doctoral
studies by Jean-Franois Cottrant and Laurent Izoret in the early 1980s, Amelie Juhin
recently related the relaxation around Cr-impurities with the colour of garnets
and spinels by coupling advanced XANES techniques and DFT calculations.
In 1980, I defended my Science Doctorate, a diploma that no longer exists
in France. A major change in my career occurred the year after, when I moved to
University of Paris 7 (now UPD) as a professor in the Earth Sciences Department.
There I met Grard Bocquier, Professor of Soil Science working on tropical soils.
As I was working on disordered materials, we found it interesting to share our
respective experiences to better understand the mineralogical and crystal-chemical controls on element transfer during weathering and soil formation. A student
interested in soil formation and mineralogy, Alain Manceau (now CNRS Director
of Research at the Grenoble Observatory), helped initiate this new approach,
beginning with the investigation of the crystal-chemical behaviour of transition
elements in the low-temperature alteration and weathering phases associated
with the Ni, Co, and Cr ore deposits, complementing the first XAFS study on
these minerals with a detailed petrological and mineralogical characterisation
(Section 18). A few years later, Alain moved to Grenoble, where the European
Synchrotron Radiation Facility (ESRF) was starting and has been successful
since. Based on Alains work we found it important to initiate a study of simpler,
synthetic systems, using the knowledge of LMCP on disordered systems. This
was the doctoral work of Jean-Marie Combes (now Executive Director of a subsidiary of Saint-Gobain Company) discussed below in Section 14. Jean-Marie was
the first student of my group to go to Stanford as a postdoc to work with Gordon,
and this stay helped Jean-Marie convert his research to XAFS investigations of
silicate melts and encouraged him to enter the glass industry.
William Bassett invited me to the 1982 American Geophysical Union (AGU)
Fall Meeting in San Francisco, which had the first special session on the applications of synchrotron radiation in the Earth sciences. There I again met Gordon
Brown, whom I first encountered at the 1979 Gordon Conference (Fig. 2.3). This
session resulted in one of the first reviews on synchrotron radiation applications
in mineralogy (Calas et al., 1984). The AGU meeting was held in the old convention centre in San Francisco. Gordon and Glenn Waychunas arranged a wonderful
day trip to Stanford University and SSRL, which was for me very informative as a
comparison with the French LURE facility. Afterward, we had a wonderful snack
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and wine party in the hills overlooking the Bay Area. It was the start of almost 30
years of collaboration on several topics of mutual interest: structure of glasses and
melts, environmental mineralogy, speciation of contaminants in soils, and energy
resources. We contributed to various review articles (Calas et al., 1987; Brown et
al., 2006a,b), including two chapters in Reviews of Mineralogy (Brown et al., 1988,
1995b), one of which was partly written in a farm courtyard in Midwestern France
between calvados breaks. It was an expedition! At a time when the Internet did
not yet exist, we had to take with us a few
kilograms of photocopies of the literature
we were reviewing, together with enough
floppy disks to avoid problems with our
early Macintosh laptop computers.
Our collaborations continued, and
in 1984 Gordon was an Invited Professor
at the University of Paris 7, bringing his
family with him. In 1997 he was awarded
the Docteur Honoris Causa degree by the
University of Paris 7. This seldomly given
degree is the highest honour in the French
University system and culminates with a
nice formal ceremony at the Sorbonne
(Fig. 2.7). In 1991-92, I went to Stanford as
Cox Visiting Professor, bringing with me
two of my three children, who came back
fluent in English. In addition, over the last
28 years, five students from LMCP have
been postdocs in Gordons group, which
have been very important periods of their
careers.
In the 1980s and 1990s, applications
of synchrotron radiation were rapidly
expanding in the Earth and environmental sciences communities, and Gordon
and I organised together or were invited
to many meetings and special sessions
devoted to this new field, including special
sessions held at the European Union of Geosciences as well as AGU and Geological Society of America (GSA) meetings, and many national and international
meetings and conferences. I vividly remember a visit by the late Joe Smith (19282007) at Stanford in 1988, for the fall AGU meeting (Fig. 2.8). Joe, who established
the Center for Advanced Radiation Sources at the APS, was enthusiastic about
forming a GeoSync subcommittee within the AGU Mineral Physics committee
to raise funding for synchrotron radiation centres.
Figure 2.7
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In 1993, I was
appointed for f ive
years to the Ministry
of H igher Education and Research,
as Advisor in Energy,
Environmental Technologies, and Mineral
Resou rces. Bei ng
relieved from some
teaching duties, I was
able to under ta ke
significant research
activity. A couple of
years later, I chaired
t h e n e w F r e n c h Figure 2.8 J.V. (Joe) Smith (centre) with Gordon Brown
(left) and Georges Calas (right) at a fence along
Nat iona l P rog ra m
the major offset along the San Andreas fault in
on soils and erosion,
the Stanford Hills resulting from the great 1906
San Francisco earthquake (photo taken in 1988).
which was the occasion to develop the
funding of molecular-scale approaches to environmental sciences in France.
Grard Bocquier had retired in the late 1980s, and we hosted at LMCP since 1988
two bright researchers from his group, Philippe Ildefonse and Jean-Pierre Muller.
This effort represented a renewal of our activity in environmental mineralogy.
Jean-Pierre Muller, Research Director at the research organisation for developing
countries (Institut de recherche pour le dveloppement, IRD) worked at LMCP
until he was appointed CEO of this organisation in 1998. He arrived with an
impressive number of laterite samples from Cameroon and Brazil, which he had
previously investigated in great detail. With Jean-Pierre and graduate students,
Thierry Allard, Blandine Clozel, and Nathalie Malengreau, we launched an investigation of the evolution of spectroscopic properties of kaolinites isolated from these
soils. Most investigated properties illustrated the interaction of kaolinites with
water outer-sphere complexes, radiation-induced defects due to the trapping
of radionuclides at the surface of clay minerals, coatings of Fe- or Ti-oxides on
clays. We adapted the EPR approach to investigate these defects and impurities
occurring at low concentrations. Thierry Allard (now CNRS Research Scientist at
IMPMC) took charge of the activities concerning the tracing of (sorbed) radionuclides in natural analogues of geological repositories of nuclear waste, with
graduate student Stphanie Sorieul (see Section 14.5), and he used spectroscopic
properties of natural colloids to obtain new information on element transport in
rivers. Philippe Ildefonse (an Associate Professor at the UPD and later Professor
at UPMC) was active in the areas of alteration processes and the structure of
poorly ordered materials such as allophanes. To determine Al-coordination in
allophanes and imogolites, we collaborated on the soft X-ray absorption spectroscopy beamline of the LURE soft X-ray facility, SuperACO, with former colleagues
from the LPS in Orsay, Anne-Marie Flank and Pierre Lagarde. Later, these data
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of nuclear glasses, I collaborated with Thierry Allard and graduate students Chlo
Fourdrin and Stphanie Sorieul on the influence of external irradiation on clay
minerals in the near field, including enhanced dissolution and amorphisation.
Since 2007 I have been a senior member of the Institut Universitaire de
France (IUF). In addition to relieving professors from a large part of their teaching
duties, this institution provides specific funding that makes research easier and
encourages the transmission of knowledge to younger colleagues. For the next
5 years, I am chairing a regional programme, supporting research on oxide materials in the Paris region. I am also involved in projects related to the exploration and exploitation of mineral resources in a sustainable environment, as well
as developing the link between materials science and cultural heritage. Some
aspects of these new studies are presented in several sections in this Perspective.
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3.
3
.1. Goldschmidts Atomistic Views of Geochemistry
I (GB) never had a chance to meet Goldschmidt as I was only 3 years old when
he died in 1947. However, excellent biographies of Goldschmidt by Mason (1992)
and Glasby (2006) provide a detailed account of the career and scientific contributions of this remarkable man who is widely acknowledged as the father of
modern geochemistry. In a lecture on March 15, 1929 to the Royal Institution
at the University of Oslo entitled The Distribution of the Chemical Elements, and
later in his classic paper presented before the Chemical Society on March 17, 1937
(Goldschmidt, 1937), Goldschmidt grouped the chemical elements into different
families (the now well-known categories lithophile, siderophile, chalcophile,
atmophile, and biophile) based on their affinities for different anions; this classification, which has stood the test of time, was based on detailed X-ray spectrographic analyses of many rocks and minerals (Goldschmidt, 1923). He also first
suggested that adsorption reactions of aqueous trace elements, particularly on
iron oxides, play a critical role in the evolution of the composition of seawater a
hypothesis that was later proven by the experimental studies of Konrad Krauskopf
(Section 4). Goldschmidts classification of the elements provides a useful guide
for predicting the affinity of ions in solution for different mineral surfaces and
natural organic matter (NOM). An example is the affinity of Type B metals such
as Hg, Ni, Cu, and Zn for reduced sulphur ligands, such as thiol groups in NOM.
In the realm of crystal chemistry, Goldschmidt is well known for his systematic
work on the crystal structures of AX, AX 2, AX 3, A 2 X 3 and other solids, from
which he derived a set of radii values of ions that bear his name (Goldschmidt
and Thomassen, 1923; Goldschmidt, 1926; Table 3.1). Goldschmidts most wellknown and lasting contribution to crystal chemistry is his radius ratio rule, with
radius ratio defined as the ionic radius of a cation divided by the ionic radius of
the anion to which the cation is bonded. This rule, when used with the ionic
radii values Goldschmidt derived, allows prediction of the most likely coordination numbers of different cations. This concept was incorporated by Pauling in
his famous rules as discussed below. However, with more modern effective ionic
radii values, systematically derived to reproduce accurately observed interatomic
distances from the sum of ionic radii (e.g., Shannon and Prewitt, 1969, 1970 and
Shannon, 1976), the radius ratio limits for various coordination numbers are not
always obeyed. For example, the radius of Si4+ coordinated by four oxygens (0.26
) falls below Goldschmidts radius ratio limit (0.225) when Si4+ is bonded to an
oxygen coordinated by Si4+ and two other cations, r( IVO2) = 1.38 . The effective
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ionic radii values derived by Shannon and Prewitt (1969, 1970) are for specific
coordination numbers, so the radius ratio principle is not applicable when using
these effective radii values (Table 3.1).
Table 3.1
Ion
Goldschmidt
Radii ()
Pauling
Radii ()
Shannon &
Prewitt Radii
()
O2
1.32
1.40
1.40
Si4+
0.39
0.41
0.40
0.26 (IV)
Al3+
0.57
0.50
0.535
Fe2+
0.83
0.75
0.78 (HS*)
Fe3+
0.67
0.645 (HS*)
Ca2+
1.06
0.99
1.00
Na+
0.98
0.95
1.02
1.32 (VIII)
K+
1.33
1.33
1.38
Mg2+
0.78
0.65
0.72
510
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own institute (The Linus Pauling Institute of Science and Medicine). A week before
my lecture on this topic, I decided to call Prof. Pauling and ask him if he would be
willing to lecture in my class on his rules as applied to minerals, as well as his
early structural studies of minerals. To my surprise, he agreed to do so and met
me at my office at 8am
on a Friday morning in
early October. I had not
yet had time to unpack
all the boxes of books
and files following my
move from Princeton to
Stanford in September,
so I stayed up most of
the night doing so
before Paulings arrival
the next morning. We
chatted for about 45
minutes prior to his
guest lecture, during
which time I told him
about my research and
Figure 3.2
X-ray apparatus at Linus Paulings desk in 1925,
had him autograph my
Gates Laboratory, Caltech.
copy of The Nature of
the Chemical Bond. Then I took him to the lecture room. He had asked that I not
advertise his lecture because he wanted to have a more intimate interaction with
the 15 or so registered students in my class. One of my teaching assistants could
not keep his mouth shut and told his friends that Linus Pauling was lecturing in my
class that morning. As a result, over 100 students and faculty colleagues showed
up for Paulings lecture in a room designed to hold about 60 people. Needless to
say, Pauling was a bit surprised to find this overflowing crowd, but it didnt affect
his presentation, which included several attempts by him to draw the structure
of several minerals on the chalkboard. He rubbed out several flawed attempts
to draw the kaolinite structure with the sleeve of his black suit coat and quickly
became covered in chalk dust. He told us a marvelous tale of his love of minerals
as he grew up in Portland, Oregon. He enrolled at age 16 in the Oregon Agricultural College (now Oregon State University) in Corvallis, graduating in 1922
with a degree in chemical engineering. During his graduate studies in the early
1920s at Caltech under advisor Prof. Roscoe G. Dickinson, he was introduced to
X-ray diffraction which he mastered and used to solve the structure of a number
of minerals and organic molecules. In addition, Pauling told the class about his
early career as an assistant professor of chemistry at Caltech during which he
solved the crystal structures of several sulphide minerals. He applied for funding
to carry out more work on metal-sulphide mineral structures with the intention
of developing a version of Paulings Rules for sulphides. However, his research
proposal was turned down, and he never got back to metal sulphides during his
long career. It is probably just as well that he did not, in hindsight, because that
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work could have distracted him from his other research pursuits in chemistry
that resulted in the Nobel Prize in Chemistry in 1954 and his work on nuclear
disarmament, which resulted in the Nobel Peace Prize in 1962.
Paulings Rules for Complex Ionic Crystals Paulings five rules are the model of
simplicity, yet they have a predictive power that is remarkably accurate. The first rule
states a coordination polyhedron of anions is formed about each cation, the cation-anion
distance being determined by the radius sum and the coordination number of the cation by the
radius ratio. The second and most important rule postulates that the state of maximum
stability of an ionic crystal is that in which the valence of each anion, with changed sign, is
equal to the sum of the strengths of the electrostatic bonds to it from adjacent cations. This
2 nd rule is also known as the Pauling electrostatic valence principle. Bond strength
(or bond valence) was defined by Pauling as the nominal charge of an ion divided by
its coordination number. Paulings third rule states cations maintain as large a separation as possible and have anions interspersed between them so as to screen their charges.
Shared edges, and particularly faces of two anion polyhedral in a crystal structure decrease its
stability. The fourth rule states in a crystal structure containing several cations, those of
high valency and small coordination number tend notto share polyhedral elements, and his
fifth rule states the number of different kinds of constituents in a crystal tends to be small.
Braggs Thoughts about Paulings Electrostatic Valence Rule Sir Lawrence Bragg
(1937) made the following analysis of Paulings second rule in his book on the Atomic
Structure of Minerals: The rule appears simple, but it is surprising what rigorous conditions
it imposes upon the geometrical configuration of a structure. In a silicate, for instance, each
silicon atom is surrounded by four oxygen atoms. These atoms have half of their valency
satisfied by the silicon, and so are left with an electrostatic charge, which is unity on our
valency scale (i.e. it is -e, the electronic charge). Aluminum within an octahedral group of
six contributes one-half to each oxygen. Magnesium or ferrous iron within an octahedral
group contributes one-third. Hence we may link a corner of a silicon tetrahedron to another
silicon tetrahedron, to two Al octahedral, or three Mg octahedra. Similarly aluminum in fourcoordination one-half, titanium in six-coordination two-thirds. Proceeding to link tetrahedra
and octahedra together in this way, we find very few alternative structures which obey Paulings law that remain open to a mineral of a given composition, and one of these alternatives
always turns out to be the actual structure of the mineral.
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digital computers. Beevers also organised the Beevers Miniature Model Unit at
the University of Edinburgh (now Miramodus Limited), which employs disabled
adults to build highly accurate, compact models of minerals and other materials.
Beevers seminar was a very memorable experience for the students and faculty
who attended because both Beevers and Pauling reminisced about the early days
of structural chemistry following Max von Laues discovery of X-ray diffraction
by crystals in 1912. I wish I had tape recorded that session!
In 1982, Jerry Gibbs and I organised a symposium entitled Applications
of Quantum Chemistry to Mineralogy at the GSA Annual Meeting, which was
held again in New Orleans, and we invited Pauling to be the keynote speaker,
and he accepted. I remember taking a limo provided by GSA to the New Orleans
airport late one afternoon to pick up Pauling. It turned out that Walter Harrison,
another of my Stanford colleagues from the Applied Physics Department, was
also invited by us to give a talk on his new valence bond theory in this symposium, and he arrived on the same flight as Pauling. Walter rode with us back to
the French Quarter, where Pauling was staying. Walter had not had the pleasure
of meeting Pauling and later told me that he was in awe during our 30 minute
drive, during which we talked with Pauling about his work on quasicrystals,
which he thought did not really exist. Our symposium began at 8am the next
morning in a room that seated about 600 people. The symposium started with
several lesser-known speakers, and we had an audience of about 50 people during
these initial talks. During the talk preceding the keynote address by Pauling,
people started pouring into the room, resulting in a standing-room-only crowd
of over 800 people just before Pauling began his talk, much to the chagrin of the
last speaker before Paulings keynote address. I introduced Pauling to enthusiastic applause from the audience, and he began his talk about the early days of
structural chemistry and his early work on the nature of the chemical bond. He
had old-fashioned lantern slides (for the younger people, look this up), including
one of an old Packard convertible automobile containing Gilbert N. Lewis, Arnold
Sommerfeld, and other famous scientists in 1925, including a young Linus Pauling
standing on the running board, so we had to rent a near-antique lantern-slide
projector for Paulings keynote address. About half way through the talk, the bulb
in the lantern-slide projector burned out, and the very large room was engulfed
in darkness. Pauling did not miss a beat in his presentation and continued on in
the dark. Fortunately, the projectionist had a spare bulb, and the lantern-slide
projector was back in operation in short order.
At the end of Paulings presentation, there were a number of questions,
followed by a mass exodus of people out of the room as I announced the next
speaker. Only about 30 people remained for that next talk. I remember the nervousness I felt when Jerry Gibbs introduced me as a speaker following the morning
break. Pauling was sitting in the front row listening to my interpretation of X-ray
photoelectron spectra of silicate minerals and glasses. At the end of my 12 minute
talk, Pauling asked me two questions about my work, which I think I answered to
his satisfaction. That evening, I organised a dinner for 8 of us, including Pauling,
Jerry Gibbs, Walter Harrison, and Marshall Newton (Brookhaven National Lab), at
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the famous French Quarter restaurant Antoines, and we were seated in the Rex
Room, where the King of Mardi Gras has his annual dinner. I sat next to Pauling
and suggested that he try the turtle soup, which he did. I also sent the first
bottle of California cabernet sauvignon back after tasting it, and was brought a
new, satisfactory bottle of the same vintage by the sommelier. Several years later
following a seminar I gave at Caltech, Barclay Kamb and his wife Linda Pauling
Kamb had a special dinner for me at their house in the Pasadena Hills, which was
originally the Pauling family home. During dinner, Linda told me that her father
enjoyed his visit to New Orleans for the GSA symposium. She also told me that
he did not like the turtle soup that I suggested he try, but he was impressed that
I sent the bottle of wine back at this fancy restaurant. I also remember another
occasion, when my late wife Nancy and I had dinner with Pauling after his wife
Ava passed away. When we arrived at his home in Portola Valley, California, he
presented Nancy with an orchid from his garden in a small chemical beaker. I
still have that beaker. He and Nancy got along well.
Linus Pauling passed away in 1994 at age 93, and the world lost a great
scientist and humanitarian I was very fortunate to know. The diversity of fields in
which Pauling made significant contributions, ranging from chemistry to mineralogy to the structure of metals to antiferromagnetism to molecular biology and
to the molecular basis of diseases (e.g., sickle-cell anemia), is truly remarkable. He
came close to solving the structure of DNA before the final solution was proposed
by Watson and Crick in 1953, which would likely have resulted in a second chemistry Nobel Prize for Pauling. In spite of his genius, Pauling was not always correct
in his thinking about the structure of solids (e.g., quasicrystals) or medicine (e.g.,
vitamin C and orthomolecular medicine). His strong opposition to the idea of
quasicrystals, which are alloys of Al with transition metals such as Fe and Cu
that have 5-fold rotational symmetry and thus consist of non-repeating units,
is an important example. I remember hearing Pauling lecture about his strong
feelings against the concept of quasicrystals in a seminar at Stanford as well as
at a meeting of the American Crystallographic Association held at Stanford. Dan
Shechtman (Israeli Institute of Technology) first discovered quasicrystals in 1982,
which were met with much skepticism by many scientists, including Pauling.
However, Shechtman was persistent in his beliefs about quasicrystals and was
finally awarded the Nobel Prize in Chemistry in 2011 for this discovery. In spite
of these shortcomings late in his career, Linus Pauling remains my scientific hero
to this day, although I have other heroes as well, particularly Jerry Gibbs, whose
contributions are discussed next.
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et al., 2008). In the early 1970s Jerry applied extended Hckel molecular orbital
(MO) theory to understand, at a more fundamental level, the inverse correlation
between Si-O bond lengths and Si-O-Si angles in framework silicates (Gibbs et
al., 1972; Gibbs et al., 1974) first recognised by Jerry and me (Brown et al., 1969).
Jerry later used more accurate ab initio Hartree-Fock-level MO calculations to
address the local bonding forces in silicate tetrahedra (Gibbs et al., 1981) that
culminated in his classic 1982 paper (Gibbs, 1982). This remarkable paper showed
that the local bonding forces controlling the bond lengths and bond angles of
H4SiO4 tetrahedral molecules and hydroxyacid molecules (H6IV T2O7: T = Si, Al,
B) in the gas phase or in an aqueous solution are very similar to those in the
tetrahedral portion of silicate, aluminosilicate, and borosilicate minerals (Gibbs,
1982). In his 2008 paper (Gibbs et al., 2008), Jerry used the concept of bond paths
developed by Richard Bader (Bader, 1991) to provide a detailed understanding
of the bonded and nonbonded electron density distributions in silicates and
sulphides and to provide a quantum mechanical basis for the empirical correlations between bond strength and bond length developed by Brown and Shannon
(1973) and others (Smith, 1953; Zachariasen, 1963; Brown and Altermatt, 1985;
Brese and OKeeffe, 1991).
In his 2008 paper, Jerry also pointed out the observation by Cremer and
Kraka (1984) that an atom in a molecule (or a crystal) is a quantum mechanical
unobservable, and he concluded that the radius of the atom is likewise an unobservable whereas the bond length is the observable. In spite of this conclusion, the
effective ionic radii of Shannon and Prewitt (1969; see also Table 3.1), because
of the way they were derived, are very useful in predicting bond lengths for
oxygen-based minerals as a function of cation and anion coordination number,
oxidation state for redox sensitive atoms, and electron spin state for first-row
transition metals.
A Quote from Richard Feynman It is instructive to begin this discussion with
a quotation used by Jerry Gibbs in his 2008 paper from American physicist Richard
Feynman, which conveys the importance of bond lengths between atoms If in some
cataclysm, all of scientific knowledge were destroyed and only one sentence was passed on
to the next generation of creatures, what statement would contain the most information in
the fewest words? I believe it is the atomic hypothesis that all things are made of atoms
little particles that move around in perpetual motion, attracting each other when they are a
little distance apart, but repelling upon being squeezed into one another. (Feynman, 1970).
Jerry and co-workers (Gibbs et al., 2003a; Gibbs et al., 2008) showed that
the correlation of average cation-oxygen distance for first- and second-row
cations, predicted from Shannon and Prewitt radii values, with their Pauling
bond strengths can be fitted to a power law expression of the following form:
R(M-O) = 1.39(s/r)0.22 where s is the empirical Pauling bond strength and r is
the row number of the M atom (Fig. 3.3a). They found that when the s values are
divided by r of the cation, the values fall along the same curve (Fig. 3.3b). This
relationship is very similar to the bond length-bond valence curves derived by
516
GEOCHEMICAL PERSPECTIVES | V O L U M E 1 , N U M B E R 4 A N D 5
Brown and Shannon (1973), Brown and Altermatt (1985), and Brese and OKeeffe
(1991) from regression analyses of experimental interatomic distances vs. bond
strengths for many oxygen-based crystal structures designed to satisfy Paulings electrostatic valence rule. Gibbs et al. (2003b, 2005) also demonstrated the
accuracy of Hartree-Fock-level quantum mechanical calculations in predicting
Si-O distances, Si-O-Si angles, and electron density distributions of the five
nonequivalent Si-O-Si linkages in coesite (Fig. 3.4). This is a remarkable demonstration of how far quantum mineralogy has progressed since the early semiempirical MO calculations on protonated disiloxy (H6IV T 2O7; T = Si, Al) groups
(Gibbs et al., 1972; Meagher et al., 1979; De Jong and Brown, 1980a,b), which
did not give accurate predictions of interatomic distances and considered only
valence electrons. However, in more accurate Hartree-Fock calculations, the
hemispherical domains in the electron localisation functions (Fig. 3.4) provide
accurate representations of bond-pair electrons along the bond vectors, and the
larger kidney-shaped domains represent lone-pair electrons on the opposite sides
of the Si-O-Si linkages. Jerry and co-workers (Gibbs et al., 2003b) also showed that
oxygens in Si-O-Si linkages in coesite with the narrowest Si-O-Si angles (the O5
oxygen bottom of Fig. 3.4) had the highest localised nonbonding electron densities (pink regions) and therefore were the most likely oxygens to be protonated.
These results offer insights about sites of electrophilic attack in silicates relevant
to dissolution reactions during chemical weathering (e.g., Xiao and Lasaga, 1994,
1996; Lasaga, 1995; Pelmenschikov et al., 2001) and in the location of hydrogen
in mantle silicates (Smyth, 1987; Smyth et al., 1991; Stebbins et al., 2009).
Figure 3.3
GEOCHEMICAL PERSPECTIVES | G O R D O N E . B R O W N J R . G E O R G E S C A L A S
517
predictions about the ways in which aqueous ions bind to reactive sites on mineral
surfaces. The bond-valence approach can also provide valuable predictions about
H-bonding to surface oxo groups at mineral-water interfaces, which cannot be
determined using current experimental methods. The constraints of Paulings
electrostatic valence principle have since been incorporated into modern surface
complexation models (Hiemstra et al., 1989a,b; Hiemstra and Van Riemsdijk,
1996; Hiemstra et al., 1996).
Figure 3.4
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4.
4.1. Langmuirs
Langmuirs concept of adsorption of molecules on solid surfaces is amazingly close to our current thinking about adsorption of adatoms and admolecules
at mineral-water interfaces, based on the results of modern spectroscopic and
X-ray scattering measurements and the applications of modern quantum chemistry and molecular modelling.
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4.2. Krauskopfs
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522
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and I watched a frustrated Dennis trying to cope with Konnies reluctance to learn
word processing on his first computer. Konnie rarely used that computer following
completion of the 3rd edition of Introduction to Geochemistry in 1995.
I had the honour of nominating Konnie for the MSA Distinguished Public
Service Award, which he won in 1994, and I served as his citationist for this
award. Part of the above description of Konnies life is from the research I did in
preparation for my citation address. What I didnt mention above is the major
public service Konnie rendered through the two books he co-authored on the
physical sciences. In 1941, Konnie published his first textbook entitled Principles
of Physical Science with co-author Arthur Beiser, a physicist. The 6th edition of
this book was published in 1974. It sold roughly 4,000 copies a year from 1941 to
1960 and was used mostly in junior colleges in courses designed for non-science
majors. Konnie teamed up with Beiser again in 1960 to produce another classic
textbook entitled The Physical Universe, now in its 12th edition (2007). This book
has been translated into Spanish and sold almost 20,000 copies in 1992 alone. It is
widely used in junior colleges throughout the U.S. I recommend that each of you
buy a copy of The Physical Universe and read it. This text is an excellent example
of how to write a book that is easily readable by a layperson but does not sacrifice
scientific accuracy and rigour. I can think of few things a distinguished scientist
can do that better serve the public interest than educating the masses about the
beauty, complexities, history, and logic of scientific discoveries.
4.3. Stumms
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present a paper at the Swiss Chemical Society Meeting in 1988 in his symposium
on Inorganic and Coordination Chemistry: Role of Surfaces. I recall that he met me
at the plane when I arrived in Zurich and handed me a fat envelope stuffed full
of Swiss francs. He told me this was spending money while I was in Switzerland.
Werner rented a special train car for the trip of his EAWAG group and all invited
speakers from Zurich to Bern. We rode to Bern in style, sipping champagne and
eating tasty Swiss food. At the symposium, I met for the first time Paul Schindler
(Professor of Chemistry at the University of Bern), who greatly impressed me with
his knowledge of coordination chemistry (Fig. 4.1a). Following the symposium,
I returned to Zurich and gave a seminar in the Chemistry Department at ETH
hosted by Werner. We became friends, and several years later (winter of 1992) I
invited Werner to Stanford as Cox Visiting Professor. While at Stanford, Werner
wrote a substantial portion of his book Chemistry of the Solid-Water Interface and
met frequently with me, my colleagues George Parks, Mike Hochella, and Jim Leckie
(former Ph.D. student of Werner and Professor of Civil & Environmental Engineering at Stanford), and our students and postdocs, including Patricia Maurice
(Professor of Civil Engineering and Geology, University of Notre Dame). When
I told Patricia I was writing a Geochemical Perspectives with Georges Calas, with
a focus on the legacy of several great 20th century scientists including Werner
Stumm, she sent me a short note about her interactions with Werner over the
years following his stay at Stanford and during a period she spent in his labs
at EAWAG. Werner was particularly supportive of young women in science like
Patricia and actively promoted their scientific careers.
Later in this Perspective (Sections 11, 12, and 16) we analyse some of the
major ideas contributed by Stumm to the field of mineral-water interface chemistry and their consistency with the results of modern experimental methods and
theory. One of Stumms major contributions was the idea that the rate at which
a mineral weathers depends on the surface charge of the mineral, which in turn
is controlled by factors such as pH and the composition of the aqueous solution.
Perspectives on the Contributions of Langmuir, Krauskopf, and Stumm
Langmuir provided us with some of the basic principles of the surface chemistry of
solids and their interactions with liquids and gases that have stood the test of time.
Krauskopf was one of the first geochemists to connect sorption reactions of metal
ions at mineral-water interfaces to the composition of natural waters, particularly the
trace element composition of the oceans. Stumm used his knowledge of coordination
chemistry and aquatic chemistry to develop models of surface complexation reactions
at mineral-water interfaces. These geochemical giants laid the foundation for our
current thinking about the chemical processes that occur at mineral-water interfaces
that will be discussed in the remainder of this Perspective.
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Figure 4.1
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525
5.
THE REVOLUTION/EVOLUTION
IN MOLECULAR-SCALE GEOCHEMISTRY
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H. Bragg (U. Leeds) and William L. Bragg (U. Cambridge) in the form of Braggs
Law, and they won the Nobel Prize in Physics in 1915 for this major contribution
that corrected some problems with von Laues formulation of X-ray diffraction and
allowed the positions of atoms in a crystal structure to be determined. William L.
Bragg using X-ray diffraction determined the atomic arrangement in a number of
minerals, including calcite, olivine, and diopside, among others (Bragg, 1937) and
showed that crystals do not, in general, consist of molecular units, as previously
thought.
The development of synchrotron radiation sources in the early 1970s
continued the revolution initiated by Rntgens discovery of X-rays by enhancing
their peak brightnesses by 25 orders of magnitude in 4th generation synchrotron
X-ray sources such as the Linac Coherent Light Source (LCLS) at the SLAC
National Accelerator Laboratory (Stanford University) (Fig. 5.1). The enormous
brightness of such a source made possible the recent structure determination of
a membrane protein nanocrystal of photosystem I using more than 3,000,000
diffraction patterns from a fully hydrated stream of nanocrystals (Chapman et al.,
2011). Even the 12 or 15 order of magnitude increase in brightness provided by
first- or second-generation synchrotron radiation sources (Fig. 5.1), respectively,
signaled the revolution that was to come in the many varied applications of
X-rays, including those in the Earth sciences. This change has greatly enhanced
the sensitivity of various X-ray methods to trace elements in complex solids and
liquids and has produced enormous decreases in the time needed for structure
determination of crystalline and amorphous solids. Synchrotron light sources
have also resulted in new X-ray spectroscopic, scattering, and imaging methods
that were not possible 20 years ago. We discuss some of these methods with
particular attention to how they have contributed to our knowledge in the area
of mineral-water interface chemistry.
What Difference Does the Intensity of a Modern Synchrotron Radiation Source
Make in Structural Studies of Solids? One of us (GB) was trained as an X-ray crystallographer. In graduate school and later I used X-ray diffraction to determine and/
or refine the crystal structures of a number of minerals, and my Ph.D. work in the
late 1960s involved refinement of the crystal structures of five olivine end-member
compositions as part of a larger project to understand the effect of cation size on the
structural chemistry of the olivines (Brown, 1980). Selection of suitable small (100
mm diameter) single crystals, their compositional characterisation, collection of X-ray
intensity data on a Picker four-circle single crystal diffractometer, and data analysis
using an IBM 370 computer took about one year. Using a modern synchrotron X-ray
source, the experimental and computational parts of my thesis work would now take
about one day using one of the protein crystallography beamlines, a modern image
plate detector, robotic crystal control, and modern digital computers.
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Figure 5.1
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GEOCHEMICAL PERSPECTIVES | V O L U M E 1 , N U M B E R 4 A N D 5
XAFS spectroscopy has become the technique of choice for in situ molecular-level
characterisation of the products of sorption reactions at solid-water interfaces.
Because of the extensive use we have made of XAFS spectroscopy in the studies
highlighted in this Geochemical Perspective issue, we need to discuss how this
method can be used to derive information on the structure and attachment
geometry of sorption complexes on mineral surfaces in contact with aqueous
solution. Before beginning this discussion, it is useful to give a brief overview of
how XAFS spectroscopy works. See Stern and Heald (1983), Calas et al. (1987),
Brown et al. (1988 and 1995a), Stern (1988), and Brown and Sturchio (2002) for a
more complete description and for references to the XAFS literature.
XAFS spectra arise from the excitation of a core-level electron of the
absorbing (or target) atom using X-rays of sufficient energy (e.g., ~7,112 eV for Fe
1s electrons). At and above this energy, the excited electron becomes a photoelectron. This energy defines the absorption edge of the absorbing element, below
which there is little interaction of the X-ray beam with the electrons except for
bound-state transitions (e.g., 1s to 3d electronic transitions in K-edge spectra of
first-row transition elements like Fe). XAFS is an element-specific spectroscopic
method that yields information on interatomic distances and types and numbers
of atoms around an absorber atom. Because of the local nature of the structural
environment that XAFS samples, there is no need for the sample to be crystalline,
and XAFS spectroscopy has been successfully applied to all types of samples,
including aqueous solutions, gases, silicate glasses and melts, fluid inclusions in
crystals, crystalline materials, and cations and anions at mineral-water interfaces.
The XAFS spectrum of an element in a particular type of sample is generated using an intense synchrotron X-ray source with a broad, continuous range of
energies. Laboratory X-ray sources simply dont provide the X-ray flux needed to
generate XAFS in reasonable time periods, particularly for low element concentrations. The X-ray beam is monochromatised in energy by Bragg diffraction from
crystals [typically Si(111) or Si(220)], and the X-ray energy is varied by changing
the monochromator crystal position. An XAFS spectrum is generated by monitoring the absorption of X-rays by a sample (or alternatively the X-ray fluorescence
yield of the sample) as a function of X-ray energy normalised by the intensity of
the incident X-ray beam. The normalised absorbance at a given X-ray energy is
given by mx = log (I0/I), where m is the absorption coefficient, x is the sample
thickness (in cm), I is the intensity of the X-ray beam after passing through the
sample, and I0 is the intensity of the incident X-ray beam. The energy region from
~30 eV to ~800 eV above the edge is referred to as the extended X-ray absorption fine structure (EXAFS) spectrum. EXAFS spectra of the metallic forms of
the first-row transition elements are shown in Figure 5.2, which also illustrates
the energy separation between the different absorption edges for this series of
elements and the element specificity of XAFS.
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529
If X-ray energy is varied only in the absorption edge region, from ~30 eV
below to ~100 eV above the absorption edge of the absorbing atom, the resulting
spectrum is referred to as the X-ray absorption near edge structure (XANES)
spectrum. Figure 5.3 shows iron K-edge XANES spectra for a number of Fe(II)and Fe(III)-bearing minerals, and the richness of fine structure that reflects the
different local environments of Fe (Wilke et al., 2001). This fine structure is caused
by interference of the ejected photoelectron wave as it travels out from the absorbing
atom and scatters off of the surrounding first-shell ligands (e.g., oxygen in the case
of the Fe-bearing oxide and silicate minerals) as well as more distant second- and
third-shell atoms surrounding the absorbing atom. The backscattered photoelectron waves travel back toward the absorbing atom and, in the process, interfere
both constructively
a nd dest r uc t ively
with the outgoing
photoelectron waves
from the absorber.
T h i s i nter ference
process, which is
highly dependent on
the number, types,
and positions of the
first, second, third,
and, in some cases,
more distant shells of
atoms surrounding
Figure 5.2
EXAFS spectra of the metallic forms of first-row the central absorbing
atom, produces a
transition elements.
modulation of the
X-ray absorption in the XANES and the EXAFS energy regions as a function of
X-ray energy, with wavelengths in the EXAFS region that are inversely proportional to the distances between the absorbing atom and its neighbouring atoms,
and amplitudes that are proportional to the atomic number of the backscattering
atom. Another effect that must be accounted for in the analysis of EXAFS data
is the phase shift that the ejected photoelectron experiences as it leaves and
re-enters the electric field of the absorbing (central) atom as well as the electric fields of the surrounding (backscattering) ligands. The extreme intensity
of synchrotron X-ray sources is needed to produce sufficient signal-to-noise
spectra in reasonable time periods (10-15min for a XANES spectrum and 30-45
minutes for an EXAFS spectrum, depending on the concentration of the element
of interest). For very dilute samples (say 50ppm of the target element), multiple
XAFS scans are required to obtain data that can be analysed to produce accurate
interatomic distances and coordination numbers.
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GEOCHEMICAL PERSPECTIVES | V O L U M E 1 , N U M B E R 4 A N D 5
Figure 5.3
Iron K-edge XANES spectra for (a) Fe(II) in sites coordinated by 4, 5, and 8
oxygens. (b) Fe(II) in sites coordinated by 6 oxygens. (c) Fe(III) in sites coordinated by 6 oxygens (after Wilke et al., 2001).
N j Fcw ( k , R )
k
g (R j )
R 2j
2 R j / l
sin 2kR j +
f ( k , R )) dR (5.1)
where S0 is the amplitude reduction factor and rf is the reduction factor for
the total central atom loss. Also, for every shell of neighbouring atoms j, Nj is
the number of backscattering atoms; |Fcw (k,R)| is the effective, curved-wave
backscattering amplitude; Rj is the average distance between the central and
backscattering atoms; g(Rj) is the (partial) radial distribution function of the
neighbouring distances around the absorbing element; l is the photoelectrons
mean free path; and Sf (k,R) is the sum of the phase-shift functions (central and
backscattering phase-shifts). This formalism is valid for any experimental XAFS
data. The phase-shift function results in a correction of the interatomic distance
of +0.2-0.7 , depending on the atom pair.
When we first began our XAFS studies in the late 1970s, there was no
adequate theory for accurate prediction of backscattering amplitudes and phaseshift functions for different atom pairs. Instead, we and other early users of
XAFS spectroscopy were required to find appropriate model compounds of
well-known crystal structure and composition similar to the unknown sample.
The XAFS spectrum of the model compound was collected and used to fit and
extract the backscattering amplitude and phase-shift functions for atom pairs of
interest. These extracted functions were then used in the analysis of the unknown
GEOCHEMICAL PERSPECTIVES | G O R D O N E . B R O W N J R . G E O R G E S C A L A S
531
compound, from which interatomic distances, coordination numbers, and DebyeWaller factors were extracted through a least-squares fitting procedure. For the
US XAFS community, this empirical approach changed dramatically in the late
1980s when John Rehr (University of Washington) and his students developed a
computer code called FEFF (for F effective) that provides accurate backscattering
amplitude and phase-shift functions for most atom pairs of interest (Rehr et
al., 1991; Ankudinov and Rehr, 2003). Rehr and his students also developed an
approach for taking into account multiple scattering of the photoelectron waves
from different atomic centres (Rehr et al., 1992; Zabinsky et al., 1995; Ankudinov
et al., 1998). It is also important to give credit to Norman Binstead and Stephen
Gurman, then at the Daresbury Synchrotron Radiation Laboratory in the UK,
who wrote the EXCURVE computer code for analysis of EXAFS data in 1982
(Binsted et al., 1982). This early software was used by many European EXAFS
workers. As a result of these advances, XAFS spectroscopy was transformed from
an art to a quantitative structural method.
When atomic vibrations are harmonic or when the distribution of interatomic distances is symmetrical, a Gaussian pair-distribution function can be
used to represent g(Rj) in equation (5.1) and can be defined for the jth shell of
neighbouring atoms as:
g (R j ) =
1
sj
exp R j R j
2 s 2j
(5.2)
where sj2 expresses the mean-square variation of Rj from the average Rj (designated Rj). In the Gaussian (or harmonic) approximation, equation (5.1) can be
rewritten as:
c ( k ) S20 rf
N j Fcw ( k , R )
k
R 2j
( 2 R /l + 2 k s ) sin 2kR + f ( k , R )
j
2
j
(5.3)
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5.2.2
With this theoretical background, lets now consider how one can use
XAFS spectroscopy to determine the mode of attachment of ions to mineral
surfaces in aqueous solutions. There are several requirements for this application of XAFS spectroscopy. First, the element of interest must not be present in
the solid (sorbent) at XAFS-detectable levels (<10 ppm), and the sorbed metal
ion (sorbate) concentration should be greater than ~50 ppm in order to produce
spectra with sufficient signal-to-noise for quantitative analysis. XAFS spectroscopy is one of the few element-specific structural methods that can be used to
characterise aqueous species and sorption reaction products of specific ions at
this low concentration level, although the concentration can be lower or higher
in practice depending on the sample matrix and synchrotron beam conditions.
Other element-specific spectroscopic methods such as NMR and Mssbauer
typically require much higher element concentrations than XAFS spectroscopy.
Electron spin resonance spectroscopy is sensitive to considerably lower element
concentrations (as low as 10 12 M in selected cases); however, its use is restricted
to S-state paramagnetic ions at room temperature, which include Cr3+, Mn2+, Fe3+,
Cu2+, Mo5+, Eu3+, and Gd3+. A second condition typically required for XAFS spectroscopy with liquid water present is that the atomic number (Z) of the absorber
should be 20 (calcium). This restriction is due to the difficulty of measuring
XAFS spectra of low-Z sorbates in the presence of water due to the low energy
of their absorption edges (<4 keV) and the fact that water strongly attenuates soft
X-rays. In practice, however, it is possible to collect near-edge X-ray absorption
fine structure (NEXAFS) spectra on elements with Z as low as 5 (nitrogen) or
6 (carbon) in the presence of thin films of water, taking advantage of the water
window (i.e. the energy range below the oxygen K-edge, 543 eV, where X-ray
absorption due to the presence of oxygen is minimal; e.g., Myneni et al., 1999;
Benzerara et al., 2004). When the sorbate is present at trace levels (<1000 ppm in
a moist sample), XAFS spectroscopy in transmission mode may be impractical,
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533
in which case X-ray fluorescence detection is used. A third condition is that the
difference in atomic number (DZ) between the sorbate metal ion and the metal
ions in the sorbent should be greater than 2. This restriction is imposed by the
similarity of electron backscattering amplitudes for elements with DZ2, which
makes them difficult or impossible to distinguish from each other, and the fact
that the fluorescent X-ray background is produced by the presence of elements
one or two Z numbers below that of the absorber. For example, both of these
problems would be encountered in a sample consisting of Co2+ sorbed on hydrous
iron or manganese oxides. In many cases, however, such as the sorption of Co2+
or Ni2+ on silica, aluminum oxides, or clays, or of Sr2+ or Pb2+ sorbed on hydrous
iron and manganese oxides, this restriction does not apply.
In the many cases where XAFS analysis of sorption products at mineralwater interfaces is feasible, it provides quantitative measures of interatomic
distances and coordination numbers for first and often second coordination shells
around specific elements (see ODay et al., 1994a) for a discussion of the accuracy and precision levels of second-shell interatomic distances and coordination
numbers derived from XAFS analysis for complex oxide structures). Second-shell
interatomic distances and second-shell coordination numbers are particularly
important for determining the way in which an adsorbate ion is bonded to a
sorbent surface (Brown, 1990), as illustrated in Figure 5.4. Additional methods,
such as transmission electron microscopy, are often essential for characterising
nano-scale precipitates containing the sorbing ion (e.g., Towle et al., 1997; Morin
et al., 2009). In the case of cation sorption to particle surfaces in the presence of
aqueous anions or organic molecules, Fourier transform infrared spectroscopy
is very useful for assessing the presence or absence of ternary surface complexes
involving aqueous metal cations and anions (or organic molecules) and mineral
surface sites (e.g., Bargar et al., 1999; Fitts et al., 1999; Ostergren et al., 2000a,b;
Ha et al., 2008).
One additional caveat about XAFS spectroscopy applied to surface
complexes involves cases where more than one type of surface complex of a given
adion is present (e.g., both inner-sphere and outer-sphere complexes), which is
likely the case in most sorption systems. In these situations, XAFS spectroscopy
sums over all species (i.e. all coordination geometries) containing the element of
interest. When one (or more) species is at a minor concentration level relative to
a dominant species (> ~50% of the total species present), XAFS is likely to detect
only the dominant species in many cases. In favourable cases, however, (i.e.
where XAFS spectral features characteristic of the major species do not interfere
with those characteristic of the minor species or when a systematic change in
the Fourier transform magnitude of second-shell features can be detected as a
function of changes in solution ionic strength), it may be possible to characterise
more than one type of complex for a specific element in a sorbate/sorbent system
using XAFS spectroscopy. For example, in the Co K-edge XAFS study of Co2+
sorption on kaolinite at Co sorption densities of 1.2 to 2.6 mol m2, ODay et al.
(1994b) were able to identify and characterise both multinuclear and mononuclear
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GEOCHEMICAL PERSPECTIVES | V O L U M E 1 , N U M B E R 4 A N D 5
Co2+ complexes. Papelis and Hayes (1996) were also able to distinguish between
inner-sphere and outer-sphere Co2+ complexes on montmorillonite as a function
of ionic strength from XAFS data.
O t her s y nc h rot ron-ba s ed
X-ray scattering methods such as
resonant anomalous X-ray reflectivity (RAXR) can detect multiple
types of adsorption complexes. This
was illustrated by Catalano et al.
(2008) who showed that both innersphere and outer-sphere arsenate
oxoanion complexes are simultaneously present at -Al2O3(0001)-water
and -Fe2O3(0001)-water interfaces.
However, this method requires
the reflecting surface of a polished
single crystal and cannot be used in
its present form on a powdered or
nanocrystalline sample, unlike XAFS
spectroscopy.
In spite of these limitations,
XAFS spectroscopy is one of the very
few methods that can provide structural and compositional information
on most types of cations and anions
sorbed at solid-water interfaces.
Moreover, it is one of the few structural methods that can, in many cases,
provide information on the speciation of selected cations and anions in
complex mixtures of phases, including
sorbed species, such as heavy metal(and metalloid)-contaminated soils
(e.g., Pickering et al., 1995; Manceau
et al., 1996, 2000; Foster et al., 1998;
Morin et al., 1999, 2001; ODay, 1999,
Ostergren et al., 1999; Roberts et al.,
2002; Everhart et al., 2006).
To illustrate how XAFS spectroscopy can be used to determine
the mode of attachment of sorbate
ions at solid-water interfaces, we next
provide an example of an XAFS spectroscopy study of environmentally
Figure 5.4
GEOCHEMICAL PERSPECTIVES | G O R D O N E . B R O W N J R . G E O R G E S C A L A S
535
Figure 5.5
In contrast, selenite uptake shows no significant ionic strength dependence, and uptake is complete well above the pH PZC of goethite, where the
particle surfaces have a strong negative charge. These observations led Hayes
and co-workers to suggest that selenite dominantly forms more strongly bound
inner-sphere complexes at the goethite-water interface under the conditions of
their experiment.
To provide more direct information about the nature of selenite and
selenate sorption complexes at the goethite-water interface, we carried out an
in situ EXAFS spectroscopy study (Hayes et al., 1987), the results and interpretation of which are shown in Figures 5.6 and 5.7. The background-subtracted and
k3-weighted EXAFS spectra (Fig. 5.6a) show that the Se K-edge EXAFS signal for
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GEOCHEMICAL PERSPECTIVES | V O L U M E 1 , N U M B E R 4 A N D 5
Figure 5.6
GEOCHEMICAL PERSPECTIVES | G O R D O N E . B R O W N J R . G E O R G E S C A L A S
537
overbonded (i.e. the bond strength sum to the oxygen would be greater than 2v.u.,
in violation of Paulings 2nd rule). However, if a proton is released from either type
of oxygen during the adsorption reaction, which would result in a reduction in
the sum of bond strengths to each oxygen of ~0.7 v.u. (Bargar et al., 1997a), and
if the VI Fe3+-O bonds lengthened slightly, then selenite, with an IIISe4+-O bond
strength of 1.3 v.u. could bond to Type A oxygens. It is less likely that the selenate
oxoanion would bond to either type of oxygen directly, even after loss of a proton
from the surface oxygens and some lengthening of the Fe3+-O bonds, because
the IVSe6+-O bond strength of 1.5 v.u. would cause both types of oxygens to be
significantly overbonded. Thus outer-sphere complexes of selenate on goethite
should be favoured. There is some evidence in support of selenate forming both
inner-sphere and outer-sphere complexes on goethite surfaces based on examination of the Fourier transform magnitude of the second-neighbour feature in
the EXAFS-derived radial distribution function as a function of solution ionic
strength (Peak and Sparks, 2002). The idea here is that increasing ionic strength
should result in increasing interference of electrolyte anions with the sorption of
weakly bound aqueous selenate oxoanions at the goethite-water interface. This
example of the use of XAFS spectroscopy to determine the binding mode of
aqueous selenium oxoanions at goethite-water interfaces is typical of the many
studies of the structure, composition, and binding mode of metal ion surface
complexes at mineral-water interfaces over the past 25 years.
Figure 5.7
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GEOCHEMICAL PERSPECTIVES | V O L U M E 1 , N U M B E R 4 A N D 5
Figure 5.8
Drawing of the goethite (-FeOOH) structure showing the two types of oxygens
(unprotonated and protonated) in the bulk structure. The goethite structure
was generated using the computer code VESTA (Momma and Izumi, 2011) and
structural data from Yang et al. (2006). The red spheres represent oxygens, and
the small light-coloured spheres represent protons. The Fe3+ ions are visible
inside the octahedra. The sizes of O, H, and Fe atoms in the drawing are not
representative of their ionic radii values.
Perspectives on the Use of Synchrotron-Based X-ray Methods to Study MineralWater Interfacial Processes Synchrotron X-ray sources and the molecular-level
methods they enable have revolutionised our views of mineral-water interface
processes, as discussed in the remainder of this Perspective. These new methods
provide detailed molecular-level information about adsorption complexes, including
how the complexes are attached to mineral surfaces. This information allows one to
determine if the complexes are inner- or outer-sphere as well as whether they are
mononuclear or multinuclear. XAFS spectroscopy is not always able to detect multiple
types of adsorption complexes that are present simultaneously at mineral-water
interfaces, and thus should be accompanied by complementary methods, some of
which, like selective chemical extractions, can help provide this information. This
type of molecular-level information on adsorption complex types is critical for understanding how strongly various contaminant and pollutant species are sequestered at
mineral-water interfaces or as separate precipitate phases, which in turn is critical
for developing robust predictive models for reactive transport of contaminants in
groundwater aquifers and soils and for predicting the potential bioavailability of these
species. Extremely intense synchrotron radiation sources, particularly the new free
electron X-ray laser sources, will revolutionise the way we study chemical processes
at solid-water interfaces, such as the making and breaking of chemical bonds during
catalytic reactions, that will lead to new understanding of the mechanisms of chemical
reactions. These sources are also likely to revolutionise our understanding of the
structure of mineral nanoparticles, particularly the structures of their surfaces, where
chemical reactions occur. The challenge for the next generation of geochemists is to
propose hypotheses about mineral-water interface reactions that can be tested using
these revolutionary new sources of photons over a broad range of wavelengths.
GEOCHEMICAL PERSPECTIVES | G O R D O N E . B R O W N J R . G E O R G E S C A L A S
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6.
Since the introduction of quantum theory in 1925 by Werner Heisenberg (Heisenberg, 1925), great strides have been made in the use of quantum mechanics and
molecular modelling (Fig. 6.1) to predict the electronic and geometric structures of matter, including minerals (Gibbs, 1982; Tossell and Vaughan, 1992;
Cygan and Kubicki, 2001). Few of these applications would have been possible
without the great advances made over the past 40 years in digital computers.
The application of modern electronic structure theory to matter has resulted in a
fundamental understanding of how atomic arrangements affect the stabilities of
molecules and solids and how these arrangements change as a function of pressure, temperature, and composition. Here we review several examples of some of
these advances relevant to interface geochemistry, in particular, the use of density
functional theory, DFT (Hohenberg and Kohn, 1964; Kohn and Sham, 1965) and
ab initio thermodynamics (Wang et al., 1998; Reuter and Scheffler, 2002; Stampfl
et al., 2002) to predict the structure of mineral surfaces and their reactivity with
water (e.g., Giordano et al., 1998; Hass et al., 1998; Trainor et al., 2004; Lo et al.,
2007; Ghose et al., 2008; Kubicki et al., 2008; Mason et al., 2010; Aboud et al., 2011;
Bandura et al., 2011a; Blanchard et al., 2012) and aqueous cations (e.g., Collins et
al., 1999; Kubicki et al., 2007; Mason et al., 2009, 2011; Zhu et al., 2009; Bandura
et al., 2011b), as well as the use of molecular modelling to predict the structure
of water (e.g., Mogelhoj et al., 2011), aqueous solutions (e.g., Driesner et al., 1998),
and mineral-water interfaces (e.g., Predota et al., 2004; Zhang et al., 2004; Vlcek
et al., 2007; Skelton et al., 2011).
Here we discuss two specific applications of molecular modelling to mineralsurface geochemistry: (a) the use of DFT to model the binding of Pb2+ ions to
the hydrated haematite and corundum
s u r f ace s ( M a s on
et al., 2009, 2011)
and (b) the use of a
number of different
ex per imenta l and
theoretical methods
to link molecularscale structure to
macroscopic properties of various ions in
Figure 6.1
Summary of molecular modelling methods used in the rutile (110)-water
molecular geochemistry.
EDL (Zhang et al.,
2004).
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541
Figure 6.2
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Figure 6.3
(upper panel) Top view of the bare TiO2 (110) showing adsorbed cations.
(middle panel) Perspective view of the TiO2 (110) surface showing H2O and
cations bonded to bridging oxygens (BO) and terminal oxygens (TO). (bottom
panel) Locations of the various cations at the TiO2 (110)-water interface (from
Zhang et al., 2004, with permission from the American Chemical Society).
GEOCHEMICAL PERSPECTIVES | G O R D O N E . B R O W N J R . G E O R G E S C A L A S
543
surface, which is consistent with a Debye length of ~100 for this ionic strength
(see equation (7.3) in section 7.1 below). In contrast, the MD simulations indicate
that water has bulk-like diffusivity and viscosity at 10-15 from the surface
plane over a wide range of ionic strengths. Zhang et al. (2004) suggested that this
discrepancy calls into question the concept of the shear plane, which may instead
be a nonequilibrium dynamical property of the dissolved ions in the vicinity
of the charged surface. This remarkable multiscale and multitechnique study
showed a general consistency between bond-valence models of surface oxygen
proton affinities and the Gouy-Chapman-Stern model of EDL structure with the
experimentally observed distribution of cations in the EDL.
Perspectives on the Utility of Theory in Providing Insights About MineralWater Interface Reactions These two examples of the application of theory (both
DFT and MD simulations) show that the combination of theory and experiment is
required to extract the maximum mechanistic information about chemical inter
actions at these interfaces, particularly the role of protons in stabilising surface sites,
which cannot be determined with current experimental methods. The challenge now
is to extend these types of studies to even more complex interfacial systems, such as
those where electron transfer reactions result in the reduction of contaminant species
(e.g., reduction of Cr6+ at the magnetite-water interface to Cr3+) or where microbial
organisms attach to mineral surfaces and influence the transformation of inorganic
contaminant species.
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7.
Over the past 40 years, surface science studies have revealed a great deal about
clean surfaces of solids under controlled conditions (Henrich and Cox, 1994;
Brown et al., 1999a; Duke and Plummer, 2002; Al-Abadleh and Grassian, 2003;
Diebold, 2003). However, clean solid surfaces exist only in ultra-high vacuum
(UHV), so the results of these studies are not generally applicable to the surfaces
of environmental solids, i.e. those in contact with environmental gases and liquids,
organic matter, fungi, or microbial biofilms. Such surfaces have structures and
reactivities that may be altered substantially by interactions with these materials,
particularly for the surfaces of redox-sensitive or anhydrous bulk solids. To handle
this complexity, we and others have employed a reductionist approach in which
experimental and theoretical studies of interfacial processes are carried out in
simplified model systems, where variables can be carefully controlled (e.g., Hayes
et al., 1987; Waychunas et al., 1993; Scheidegger et al., 1997; Towle et al., 1997;
Thompson et al., 1999; Sherman and Randall, 2003). This is followed by studies
of increasingly complex model systems, using both UHV and in situ methods,
ultimately approaching the complexity of natural systems (e.g., Templeton et al.,
2001; Fandeur et al., 2009; Ona-Nguema et al., 2009; Ha et al., 2009, 2010; Lee et al.,
2010; Wang et al., submitted-a). In order to assure that the model systems chosen
for study are relevant to the naturally occurring phenomena we wish to understand, parallel field studies of real environmental samples (e.g., soils polluted
with arsenic or lead) are also essential (e.g., Foster et al., 1998; Morin et al., 1999;
Cancs et al., 2005). To illustrate this approach, we present a number of examples
of experimental and theoretical studies of mineral-water interface reactions in
model systems of increasing complexity as well as in real environmental systems.
We begin our discussion of mineral-water interface chemistry with a brief
review of the acid-base chemistry of mineral surfaces, in contact with water,
followed by a summary of current knowledge of the EDL, which extends 10 to
20 from the solid surface into solution. Ion concentrations in, and properties
of, the EDL differ substantially from those of the bulk aqueous solution because
of the pH-dependent charge developed by the solid surface. This region contains
ions and molecules of various types, including those with charge opposite to that
of the solid surface.
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545
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GEOCHEMICAL PERSPECTIVES | V O L U M E 1 , N U M B E R 4 A N D 5
theory and Paulings 2 nd rule, i.e. that bridging oxygens that satisfy Paulings
2nd rule should be less reactive than non-bridging oxygens except at pH values
well above the pH PZC of the sorbent, where surface protons are weakly bound. A
similar bond-valence argument can be made to explain the relatively low reactivity of the silanol surface of clay minerals such as kaolinite (Sposito, 1984). In
contrast, the surfaces of iron and manganese oxides and the edges of aluminol
layers of clay minerals should be more reactive because a significant fraction of
the surface oxygens are underbonded in the Pauling sense, which is corrected
prior to adion adsorption by additional bonding of these oxygens to exchangeable hydrogen ions.
A more recent ATR-FTIR study of water-exposed powdered quartz surfaces
by Koretsky et al. (1997) found evidence for only one surface OHstretching band
at 3745 cm1 following heating of the sample to 600 C for 24 hours, although a
number of other weaker bands in the 3500 to 3300 cm1 region were observed and
assigned to internal OH groups associated with Li, Na, and other impurity ions
proposed as defects in the quartz structure. The single sharp, slightly asymmetric
band at 3745 cm1 was assigned to terminal silanol groups (Si-OH), although the
possibility of geminal (Si(OH)2) groups was also discussed in light of the peak
asymmetry mentioned earlier. Evidence for geminal (Si(OH)2) groups has been
provided by an earlier proton NMR study of powdered quartz (Bergna, 1994).
The results of many other surface-sensitive spectroscopic studies of the
interaction of water with a number of metal-oxide surfaces are summarised
in Thiel and Madey (1987) and Noguera (1996). Some of these studies report
temperature programmed desorption (TPD) measurements of molecular water
or hydroxyl release from clean metal-oxide surfaces in ultra-high vacuum environments (e.g., Stirniman et al., 1996). Such TPD studies typically involve the
adsorption of water (or D2O) at low temperature (e.g., 100 oK), where ice rather
than liquid water is stable, followed by monitoring desorption of hydroxyl or
molecular water with increasing temperature. Thus, such studies are not particularly relevant to the interaction of liquid water with mineral surfaces. A more
direct approach to studying the interaction of water vapour and liquid water with
mineral surfaces is discussed in Section 9.
We have cited evidence above that the surfaces of metal-(oxyhydr)oxides
and simple silicates contain ionisable hydroxide sites, XOH. Protonation and
deprotonation of XOH, in turn, result in pH-dependent surface charge and
potential both zero at a solid-specific pH, the pH PZC (Parks, 1965, 1967, 1990;
Sverjensky, 1994).
GEOCHEMICAL PERSPECTIVES | G O R D O N E . B R O W N J R . G E O R G E S C A L A S
547
colloidal particles in contact with water develop a surface charge. This concept
was followed by the idea that the collective charge of counter-ions in the aqueous
solution surrounding a solid particle counteracts this surface charge (Gouy, 1910;
Chapman, 1913). This early work led to the concept of an EDL at solid-water
interfaces (Fig. 7.1) in which counter-ions (ions with a charge opposite to that
of the solid surface) and co-ions (ions with a charge of the same sign as that of
the surface) are diffusely distributed in the EDL, with counter-ions increasing in
concentration toward the solid surface and co-ions increasing in concentration
away from the solid surface, until some steady state concentration is reached at
10-20 from the interface (i.e. the bulk aqueous solution).
Stern (1924) added the idea that the charge in the EDL is separated from
the surface charge because of the finite size of counter-ions, which limits their
approach to the surface. The so-called Stern layer represents this region just
adjacent to the solid surface
with no counter-ions. Stern
and later Grahame (1947)
added the concept that
aqueous ions can specifically
adsorb at the Stern (or 0)
plane. More modern versions
of the EDL are shown in
Figure 7.2. The 0-plane in
Figure 7.2a corresponds to
the Stern or compact layer.
The surface charge
of the solid influences all
charged and polar species
Evolving models of the EDL formed by
Figure 7.1
in solution. Ions of opposite
an electrolyte solution in contact with a
negatively charged surface. Shown below charge (counter-ions) are
each model is the distance (z) dependence attracted toward the surface,
of the electrostatic potential (F) (after where some may come
Bedzyk et al., 1988).
close enough to touch the
surface. However, unless
a chemical bond forms, these ions are indifferent to the surface, remain fully
hydrated, and accumulate or adsorb in a diffuse layer. They are attracted to the
surface by its electrical potential, but are randomised by thermal kinetic energy
so that their concentration is likely highest at a distance approximately equal to
the fully hydrated ionic radius of the counterion, and decays exponentially with
distance into the solution until equal to the bulk solution concentration (Fig. 7.2a).
Similarly, indifferent ions of like charge (co-ions) are repelled from the surface.
The net result is at least two layers of charge the net charge carried by the
positively adsorbed counter-ions and negatively adsorbed co-ions, and the charge
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549
Figure 7.2
(a) Schematic view of the EDL at a negatively charged mineral-aqueous solution interface showing complexes and the various
planes of the EDL model (after Brown and Parks, 2001). (b) Computer simulation of the goethite-water interface (from Zarzycki
et al., 2010, with permission from the Croatian Chemical Society).
on the surface itself resulting in an EDL. The Helmholtz-Gouy-Chapman-SternGrahame EDL model provides for the possibility that some ions appear to bind
specifically or more strongly than simply electrostatically, and thus approach the
surface more closely. This model also introduced a specific adsorption potential,
which can be expressed (after Hunter, 1987) as a contribution to the Gibbs free
energy of adsorption. Thus for each counter-ion or co-ion, i,
DGads = zieyi + DGchem + DGsolvation
(7.1)
where zieyi is the free energy associated with the long-range Coulombic or electrostatic interaction at the location of the adion, DGchem is the free energy associated with formation of a chemical bond, i.e. a covalent bond or a hydrogen bond,
etc., between the adion and the surface, and DGsolvation is the free energy of solvation, which is a function of zi2e2/r i[1/eint 1/e bulk water], where zi is the charge on
the solvated ion, e is the electron charge, r i is the radius of the solvated ion, eint is
the dielectric constant of water in the interfacial region of the EDL, and ebulk water is
the dielectric constant of bulk water (James and Healy, 1972c). DGsolvation is positive because ebulk water is greater than eint (Fig. 7.2a), which means that Men+-H 2O
bonds (where Men+ represents a metal ion) must be broken in order for Men+ to
form a direct chemical bond to surface functional groups.
If DGchem is nil or if DGsolvation >DGchem , then the adion is indifferent,
adsorption is non-specific, and the ion is in the diffuse layer, beyond the outer
Helmholtz plane (OHP, Fig. 7.2a). If chemical bonding takes place, DGchem is
finite and the adion is specifically adsorbed as an inner- or outer-sphere (or
ion pair) surface complex, located at one of the inner Helmholtz planes (IHPs
in Fig. 7.2a). Some surface complexation codes (e.g., HYDRAQL, Papelis et al.,
1988) simplify this model, locating inner-sphere complexes on the surface, thus
collapsing that portion of the IHP into the surface (the 0-plane, distance from the
surface = 0), where the potential is y0 (Fig. 7.3). Outer-sphere adion complexes
that are specifically adsorbed are located in an IHP designated IHP2 (Fig. 7.2a).
Supporting electrolyte ions, if modelled as ion-pair surface complexes, are located
in the IHP2 plane as well. The diffuse layer begins at the OHP, in keeping with
Grahames original definition.
Observable surface potentials are reduced by adsorption of indifferent
ions, but cannot be reversed, because the only attraction is electrostatic. Because
DGchem is independent of Coulombic interactions, specifically adsorbed ions can
adsorb on surfaces of like charge, and thus can reverse surface charge (Fuerstenau,
1970; James et al., 1977). James and Healy (1972b) and Hunter (1987) discuss
these criteria and review data that lead to the following generalisations. An
electrolyte or ion that adsorbs specifically causes a shift in the pH point of zero
charge (pH PZC), causes asymmetry in the potentiometric titration of a suspension, reverses surface charge or potential, or adsorbs on a surface of like charge.
Ignoring direct adsorption measurements, electrochemical criteria alone yield
the following typical results. Sodium and potassium nitrates and perchlorates
are indifferent with respect to oxides and silicates, whereas chloride is indifferent
550
GEOCHEMICAL PERSPECTIVES | V O L U M E 1 , N U M B E R 4 A N D 5
with respect to Al2O3 but specific with respect to Fe2O3. Ba 2+ and sulphate are
specifically adsorbed by alumina, although sulphate appears indifferent to the
negative surface. The alkaline earth cations are specifically adsorbed by SiO2,
MnO2, TiO2, and some Fe2O3 preparations with the following order of affinities: Ba 2+>Sr2+>Ca 2+>Mg2+ (Fuerstenau et al., 1981; Jang and Fuerstenau, 1986).
On Al2O3 surfaces and some Fe2O3 preparations, the alkaline earth cations are
specifically adsorbed but with relative affinities that are the reverse of those for
TiO2. This reversal in sequence has been explained in terms of the structure of
water at the interface and the structure-making or structure-breaking properties
of the adion (Brub and DeBruyn, 1968; Huang and Stumm, 1973). Co2+ ions
behave specifically and precipitate on TiO2. In addition, Co2+ ions are specifically
adsorbed on g-Al2O3, and Mo(VI) oxoanions are indifferent on silica but specific
on alumina (De Boer et al., 1993).
The charge on a metal-(oxyhydr)oxide surface, and thus the surface potential, depends on the degree of protonation of surface hydroxyls and the extent of
adsorption of charged inner-sphere species. The stronger the average Bronstead
acid character of the surface hydroxyls, the lower the pH PZC, which also depends
on the structure, composition, and bonding in the solid substrate. The potential
at the OHP is related to the surface potential and surface charge through a
capacitance term, which is a function of surface to OHP distance and a local
dielectric constant (Hunter, 1987). In the diffuse layer, ions are expected to follow
a Boltzmann distribution in which the potential field decays away from the OHP
in a near-exponential fashion according to the Gouy-Chapman theory (Carnie
and Torrie, 1984; Sposito, 1984; Hunter, 1987) as described by
(7.2)
where yx is the potential (in volts) at distance x (in m) from the OHP, yd is the
potential at the OHP, z is the valence of the counter-ion, e is the charge on the
electron (= 1.602 10 19 Coulombs, C), k is Boltzmanns constant (1.38 10 23 JK1),
T is temperature (in Kelvins), and k is the reciprocal of the Gouy layer thickness
(sometimes referred to as the Debye-Huckel parameter; in m1), which at 25 C
in water is given by
k = {(2000 F2 Si zi2ci) / (ere0 RT)} 1/2
(7.3)
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Hayes and Leckie (1986) presented a modified version of the TLM that allows
both inner- and outer-sphere
complexes to be considered
explicitly, resulting in an
improved treatment of ionic
strength effects (Fig. 7.3).
This is achieved, in part, by
allowing a significant fraction
of the counterions traditionally placed in the diffuse layer
to associate with surface sites
as outer-sphere complexes
in the b-plane (or IHP 2 in
Fig. 7.2a). Empirically, these
adions are weakly bound relative to those that bind directly
to the surface and have been
termed non-specific. The
association with specific sites
and the fact that the sorption
equilibrium constant has a
finite value over and above
purely electrostatic contribuFigure 7.3
Schematic view of the triple layer model tions, however, places them
of the mineral-water interface showing
within Grahames (1947) defipossible inner-sphere (Pb2+) and outer
+
2
sphere (Cl , Na , SO 4 ) complexes and nition of specifically adsorbed
the drop off of the electrical potential ions.
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Leckie, 1987; Hayes et al., 1988; Parks, 1990), direct determination of the structure
and nature of surface complexes and the structure of the EDL is not possible by
these methods alone (Westall and Hohl, 1980; Brown, 1990).
One of the experimental methods for determining EDL structure is
synchrotron-based X-ray standing wave fluorescent yield (XSW-FY) spectroscopy. This method was introduced by Batterman and Golovchenko and co-workers
in several classic papers (Batterman, 1964, 1969; Golovchenko et al., 1974; Cowan
et al., 1980) that pointed out the utility of both long-period and short-period (or
Bragg) XSW-FY spectroscopy for probing the distribution of metal ions at buried
interfaces. Subsequently, the long-period XSW-FY (LP-XSW-FY) method has
been used to probe the vertical distribution of atoms at a number of different
types of interfaces, including electrochemical interfaces (Bedzyk et al., 1986;
Abruna et al., 1990); biological membranes (Bedzyk et al., 1988, 1989, 1990; Wang
et al., 1991, 2001), organic thin films (Bedzyk, 1992), mineral-water interfaces
(Trainor et al., 2002a; Levard et al., 2011b) and mineral surfaces coated with
microbial biofilms or thin organic films (Templeton et al., 1999, 2001, 2003b; Yoon
et al., 2005a). Bedzyk and Cheng (2002) and Trainor et al. (2006) provide excellent
reviews of these various applications of LP-XSW-FY spectroscopy through 2006.
The first application of the LP-XSW-FY method to ion distributions at solidwater interfaces was carried out by Mike Bedzyk and co-workers (Bedzyk et al.,
1988), who probed the distribution of Zn2+ ions at a membrane-coated solid-water
interface. The X-ray standing wave field generated by reflection of X-rays off of
a reflecting surface is shown schematically in Figure 7.4. The multilayer solid
sample used in this experiment is described and illustrated in Figure 7.5, which
also shows the experimental (solid circles) and theoretical (solid lines) Zn Ka
fluorescent yield data for this sample at different pH values (2.0, 4.4, 6.8) as well
as the X-ray reflectivity profile as a function of X-ray incidence angle q (in mrad).
The results of this study showed that the Helmholtz model of the interface region
(Zn2+ ions only at the solid-water interface), as illustrated by the dashed line in
Figure 7.5, panel b, is incorrect. They also showed that the distribution of Zn2+
ions in the interfacial region is not constant. Instead, the results are consistent
with a distribution of Zn2+ ions predicted by the Gouy-Chapman EDL model.
Paul Fenter and co-workers (Fenter et al., 2000a) carried out one of the first
small-period (Bragg) XSW-FY spectroscopy and reflectivity studies of the EDL
at a mineral-water interface. They examined the locations of Rb+ and Sr2+ ions at
the rutile (110)-solution interface at pH values ranging from 7.9 to 10.9, which is
well above the pH PZC of rutile (5.5). More specifically, they determined the Rb+
and Sr2+ locations in the compact (or Stern) layer as well as the partitioning of
these ions between the compact and diffuse layers under in situ conditions using
XSW triangulation methods that involve XSW generation from several different
lattice planes. The heights of Rb+ and Sr2+ above the rutile (110) surface in the
compact layer were found to be at 3.35 and 2.75 , respectively, in this initial
study. The fact that the difference in height (0.60 ) is greater than the difference
in their ionic radii (D = 0.46 : Shannon, 1976) suggests that their positions are in
part a function of ionic charge and adsorption geometries and not exclusively the
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GEOCHEMICAL PERSPECTIVES | V O L U M E 1 , N U M B E R 4 A N D 5
size difference. By combining the XSW-derived z-positions of Sr2+ ions with the
ex situ EXAFS-derived average Sr-O bond length (2.60 ), Fenter and co-workers
concluded that Sr2+ ions in the compact layer are bonded in an inner-sphere,
tetradentate fashion to two adjacent non-bridging oxygens above Ti4+ ions in
the surface and two oxygens bridging between two Ti4+ ions. The maximum
partition coefficients of Rb+ and Sr2+ between the diffuse and compact layers were
estimated to be 0.5 0.25 for Sr2+ and 0.7 0.3 for Rb+.
Figure 7.4
Schematic drawing of the X-ray standing wave field formed by the interference between the incident, E0, and specular-reflected, ER , plane waves above
a reflecting surface. The antinodes of the standing waves are parallel to the
surface and have a period of D = l/2sinq. Because total external reflection
of the wavelength X-rays occurs at incident angles, q, that are less than
0.5, D typically varies between 100 and 1000 (after Bedzyk et al., 1988).
These results suggest that Sr2+ is partitioned about equally between the
compact and diffuse layers, and that about 70% of the Rb+ ions occur in the
compact layer. However, the Kads for Rb+ is about 150 times smaller than for Sr2+.
The XSW results for Sr2+ and Rb+ are not readily comparable because of the lower
pH values used in Rb+ experiments (7.9 to 9.9) vs. the Sr2+ experiments (10.6-10.9)
as well as the slightly higher ionic strengths in the Rb+ experiments.
This early study was followed by the more detailed multi-technique study
of Zhang et al. (2004) described above. The heights of Rb+ and Sr2+ above the
unrelaxed Ti-O surface plane were found to be 3.44 and 3.12 , respectively,
differing somewhat from the earlier study by Fenter et al. (2000a). All sorbed
species were found to form inner-sphere complexes, with the specific binding
geometries of the cations and their reaction stoichiometries dependent on
the ionic radii of the cations (Fig. 7.6). This finding is consistent with Dimitri
GEOCHEMICAL PERSPECTIVES | G O R D O N E . B R O W N J R . G E O R G E S C A L A S
555
Sverjenskys hypothesis that surface oxygens of metal oxides with high dielectric
constants are able to hydrate sorbed cations at least as effectively as solvent
molecules (Sverjensky, 2001).
Figure 7.5
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surface charge should cause water molecules near the surface (estimated to be
at least three monolayers) to be oriented with their positively charged hydrogen
dipoles pointing toward the surface, leading to a high degree of order in these
layers relative to bulk water. At pH 2, where the surface charge of quartz is neutral
or slightly positive, a considerably smaller but still relatively strong SFG signal
was measured indicating a relatively high degree of order of interfacial water
molecules and leading to the suggestion that the negative oxygen dipoles of these
water molecules point toward the quartz surface. At pH values between 3 and 12,
the SFG signal was found to be very weak, indicating that water molecules are
disordered in the vicinity of the quartz surface over this pH range. This conclusion
is consistent with the model for water at the quartz-water interface developed by
Schlegel et al. (2002).
High-resolution X-ray reflectivity studies of mineral-water interfaces are
consistent with adsorbed water density being equivalent to the site densities of
exposed metal cations at the surfaces of calcite(104) (Fenter et al., 2000b) and
barite(001) and (210) (Fenter et al., 2001), which are lower than that of bulk water.
The fully hydrated quartz (100) and (101) surfaces have adsorbed water densities
consistent with the surface-site density of silanols (Schlegel et al., 2002), and
the orthoclase (001) surface (Fenter et al., 2000c) has adsorbed water density
consistent with the surface site density of silanols plus aluminols. In all of these
cases, there is little or no observable perturbation of water above the surface,
and relaxations, which may affect the outermost several unit cells, are limited
to no more than a few tenths of an ngstrom at most. At the orthoclase (001)
surface, K+ ions are absent, presumably exchanged by hydronium ions (Fenter
et al., 2000c). As discussed above, the muscovite(001)-water interface is unique
in having two adsorbed water layers forming a dense water layer, plus additional
water structuring as far as 10 from the surface (Cheng et al., 2001).
Perspectives on the Nature of Solid-Water Interfaces We now have the means
to determine some of the distances between mineral surfaces and specifically and
non-specifically adsorbed ions in the EDL under in situ conditions (i.e. in the presence
of bulk water at ambient T and P). The various synchrotron-based molecular-scale
probes described above can provide this information. In addition, MD simulations of
the solid-water interface provide strong support for the EDL model as conceptualised
by Helmholtz, Gouy, Chapman, Stern, and Grahame and as used to predict the properties
of colloidal particles. Some of the studies discussed above point to structured water in
the layer adjacent to the mineral surface and, by inference, differences in H-bonding
and electrostatic interactions in the surface water layer relative to bulk water. The first
monolayer of water at the quartz-, calcite-, barite-, and orthoclase-water interfaces
was found to be less dense than bulk water, whereas water at the muscovite-solution
interface was found to be about the same density as bulk water. Based on these results,
there does not appear to be a simple correlation of the density of the first monolayer
of water with the hydrophobicity or hydrophylicity of the mineral surfaces as one
might expect. These differences between interfacial and bulk water undoubtedly
affect the interactions of metal ions with hydrated mineral surfaces. It is likely that
the dielectric properties of the mineral affect these interactions as well as influence
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the hydration of interface ions, with high dielectric constant mineral surfaces affecting
the hydration sphere of cations within the -plane more than low dielectric constant
mineral surfaces. Thus high dielectric phases like TiO2, Fe3O4, and MnO2 (Table 10.1)
are expected to favour loss of waters of hydration around cations in the -plane of
the EDL, as the work required to remove waters of solvation should be smaller the
higher the dielectric constant of the sorbent. In contrast, lower dielectric phases like
alumina and quartz are expected to favour fully hydrated cations. This reasoning is
also consistent with the prediction of solvated alkali and alkaline earth cations in the
inner-layer region of metal-(oxyhydr)oxides with lower dielectric constants.
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8.
Water is essential for life and comprises ~57% of the total body weight of an
average human adult (Guyton and Hall, 2011). It also covers ~70% of Earths
surface in the form of the oceans. Water is one of the strangest liquids known,
with many anomalous properties relative to normal liquids such as its increased
density upon melting, its density maximum at 4 C, and its abnormally high
surface tension (Chaplin, 2012). As the main liquid in Earths critical zone, it
is also the major player in low-temperature geochemical processes, including
geochemical processes that occur at mineral-water interfaces. Given its importance in these and many other areas, it is reasonable to assume that we have
a relatively clear picture of the atomic-level structure of water. Most inorganic
chemistry textbooks show the structure of a water molecule as consisting of two
O-H bonds ~0.96 in length with an H-O-H angle of ~104.5. Two lone pairs
of electrons have been proposed to reside on the opposite side of the oxygen
atom from the two O-H bonds, resulting in a significant dipole moment for the
water molecule and lone pair-lone pair and lone pair-bond pair repulsions that
give rise to an H-O-H angle less than the ideal tetrahedral angle of 109.5. This
simple picture of the isolated water molecule has been modified more recently
based on the results of ab initio quantum chemical calculations, which favour less
pronounced electron density where the lone pairs are purported to reside, and an
sp2 hybrid model of oxygen, with an unhybridised pz oxygen orbital (Laing, 1987).
H-bonding of each water molecule to four adjacent water molecules through two
acceptor and two donor bonds is thought to result in a 3D continuous random
network of water molecules, with the O-H distances slightly longer (~1.1 )
than the isolated water molecule because of H-bonding, which weakens the
covalent bonding and reduces the repulsion between electron orbitals (Soper and
Benmore, 2008). This is the conventional picture of the static water molecule and
liquid water with which many of us are familiar. Is this the real story, however?
Periodically, this conventional model of water is challenged based on new
observations. One of the most famous of these challenges occurred in the 1960s
when polywater was first reported by Deryagin and co-workers (see Deryagin,
1970 for a review). It was claimed that polywater had a lower freezing point
and a higher boiling point as well as a density that was 20% greater than that of
normal water. As discussed in Clark et al. (2010a), strong differences of opinion
about the existence of polywater were voiced in numerous publications by
various groups (e.g., Lippencott et al., 1969; Willis et al., 1969; Bascom et al., 1970;
Hildebrand, 1970; Rosseau, 1971). Leland Allen, a quantum chemist I (GB) got to
know when I was at Princeton in the early 1970s, provided theoretical proof
that polywater existed (Allen and Kollman, 1970). However, when experimental
evidence against the existence of polywater was published, Allen and Kollman
(1971) published a second paper showing theoretical evidence against polywater.
There was even concern by the public that if polywater came into contact with
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normal water supplies, it would cause polymerisation of normal water and lead
to a doomsday scenario (Time Magazine, 1969). This idea is similar to the one
put forth in Kurt Vonneguts science fiction novel Cats Cradle (Vonnegut, 1960).
This controversial idea was put to rest in 1970 based on X-ray photoelectron and
IR spectroscopy studies, which strongly suggested that polywater was simply
normal water containing a salt similar to human sweat (Rousseau and Porto,
1970; Barnes et al., 1971; Davis et al., 1971).
My Stanford colleague Anders Nilsson and his Stockholm University
collaborator Lars G. M. Pettersson, took a scientific journey that led them into
uncharted waters and resulted in a new mixture model of the structure of
liquid water that has proven to be very intriguing but quite controversial. Water
structure models have traditionally fallen into two categories mixture models
and continuum models. It is interesting to note that one of the first mixture
models for water was proposed by Rntgen (1892), who would discover X-rays
three years later. Unlike polywater, however, this new model has a very strong
scientific basis and poses a serious challenge to the conventional wisdom about
water. We would like to share this story because of its potential impact on our
thinking about the most universal solvent on Earth and because of the model
this scientific saga provides for the scientific method. One of us (GB) has followed
this interesting story since its beginning when Anders and Satish Myneni (now
Professor at Princeton) were doing NEXAFS experiments together at the ALS
using a UHV chamber developed by Satish. According to Satish, he and Anders
were having lunch at the LBNL cafeteria and were discussing the possibility
of measuring the oxygen K-edge near edge X-ray absorption fine structure
(NEXAFS) spectrum of water that afternoon. The main problem was that they
had no suitable sample cell for water that was compatible with the UHV chamber.
On the way out of the cafeteria, Satish picked up a plastic straw and took it back
to the beamline. His idea was to cut a small section from the straw, immerse
it in water, then attach the water-filled straw to the sample flange in the UHV
chamber. Satish and Anders did so and found that the surface tension of water
was sufficient to hold the water in the plastic straw even in the UHV system.
They carried out a normal-incidence excitation/grazing-angle detection NEXAFS
experiment on liquid water at the oxygen K-edge (532eV) (Myneni et al., 2002)
and were surprised by what they saw.
Figure 8.1a presents a later version of the oxygen K-edge NEXAFS spectrum of liquid water, with no self absorption, compared with the spectra of gasphase water and ice, showing that the electronic transitions in the condensed
phases are significantly different from those in the gas phase (Nilsson et al., 2010).
Anders and his co-workers attributed the disappearance of the sharp peaks in the
spectra of the condensed phases to the overlap of unoccupied molecular orbitals
on neighbouring molecules, which leads to a major rehybridisation. The oxygen
K-edge NEXAFS spectrum shows a stronger pre-edge feature relative to the spectrum of ice, but a much weaker post-edge than for ice. The main edge is present
in both spectra but is much more intense for liquid water. Figure 8.1b shows
the origin of the oxygen K-edge NEXAFS of water, which is due to electronic
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Figure 8.1
(a) Oxygen K-edge NEXAFS spectra of ice, liquid water, and water vapour
(after Nilsson et al., 2010). (b) Schematic molecular orbital diagram showing
the origins of the NEXAFS and XES of liquid water. The 1b2 , 3a1, and 1b1
orbitals are occupied, whereas the 4a1 and 2b2 orbitals are unoccupied. Note
the large difference in the spatial extent of the unoccupied vs. the occupied
orbitals, indicating that occupied orbitals are in closer proximity to the water
molecule (from Nilsson and Pettersson, 2011, with permission from Elsevier).
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Since publication of the Wernet et al. (2004) paper, there have been a number
of studies that have attempted to refute or reinterpret the Nilsson-Pettersson
model for liquid water, mostly from groups that do neutron and X-ray scattering
(e.g., Soper, 2005, 2007, 2008; Clark et al., 2010a,b; Soper, 2010, 2011; Soper et
al., 2010) or MD simulations of liquid water (e.g.,
Hura et al., 2003; Horn et
al., 2004). The classical
and ab initio MD simulations almost uniformly
support the conventional
cont i nuou s ra ndomnetwork model of liquid
water because of the
Figure 8.2
(a) Schematic illustrations of single-donor interatomic potentials
water species and (b) double-donor water they use. However, ab
species (after Wernet et al., 2004).
initio MD simulations of
liquid water using a new
water potential, which explicitly includes van der Waals interactions derived
from density functional theory, support the high-density-like water structure
rather than the mainly tetrahedral continuous-network model (Mogelhoj et al.,
2011). Over the past few years, Nilsson and co-workers have presented various
new evidence supporting their new model. These include new small-angle X-ray
scattering data suggesting an inhomogeneous mixture of different domains (see
Fig. 8.3; Huang et al., 2009), new X-ray Raman-based XAFS data (Bergmann et al.,
2007; Nilsson et al., 2010); new XES data supporting two structural motifs in liquid
Figure 8.3
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water (Tokushima et al., 2008, 2010), or new interpretations of X-ray and neutron
scattering data for water (Leetma et al., 2008; Wikfeldt et al., 2009, 2010). However,
the opposition usually follows with another paper that attempts to reinterpret the
new results in a way favourable to the conventional continuous random-network
model (e.g., Clark et al., 2010a,b; Soper et al., 2010), which is followed, in turn, by
another paper by the proponents of the new model.
The interested reader is referred to two recent reviews of the evidence
in favour of the new model (Nilsson and Pettersson, 2011) and the opposition
view (Clark et al., 2010a) and also to a recent paper (Guo and Luo, 2010) that
reports new X-ray emission spectroscopy data on liquid water that supports the
disordered-mixture model of water. There has even been an attempt to refute
the Nilsson-Pettersson model of liquid water using a modified Pauling bondvalence approach (Bickmore et al., 2009); however, one must keep in mind that
the bond-valence approach does not account for directional bonds resulting from
hybridisation of atomic orbitals, which are important in water molecules.
Perspectives on Our Understanding of What Water Is It has been interesting
to follow this controversy, which serves as a useful model of the scientific method.
Conventional wisdom is challenged by a new model based on new observations and
new theoretical interpretations. Others, who favour the conventional wisdom, counter
with alternative explanations and new measurements of their own in an attempt to
cast doubt on and refute the new model. And so it goes in this interesting case. We
do not know which model is correct at this point in time, although one of us (GB)
favours the Nilsson-Pettersson model after careful examination of the evidence. Water
remains an enigma, as captured in a quotation from D. H. Lawrence (1929):
Water is H 2O, hydrogen two parts, and oxygen one
But there is also a third thing that makes it water
And no one knows what that is
(I believe God knows)
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9.
9.1. Introduction
Arguably the most fundamental chemical reaction involving mineral surfaces in
Earth-surface environments is their interaction with aqueous solutions (Stumm
et al., 1987; Brown, 2001). Even in air, metal-(oxyhydr)oxide surfaces are likely to
have multiple monolayers (ML) of sorbed water. For example, -Al2O3 surfaces
have been shown by thermo-gravimetric analysis to have the equivalent of one
monolayer of water at a relative humidity (RH) of ~35% and the equivalent of more
than 20 ML at 95% RH (Yan et al., 1987). In the case of high surface area silica,
Miyata (1968) found the equivalent of 26 monolayers of water at 98% RH, and
Pashley and Kitchener (1979) found the equivalent of 50 ML of water on quartz
at 100% RH. When exposed to liquid water, many metal-oxide surfaces become
hydrated or hydroxylated over time. The outer-most surfaces of aluminum oxides
such as -Al2O3, for example, hydroxylate rapidly (within a matter of minutes;
Liu et al., 1998d) followed by hydroxylation of more extensive regions, resulting in
conversion of the surface region of aluminum oxides into boehmite (a-AlOOH)
and bayerite (g-Al(OH)3; Dyer et al., 1993; Lee and Condrate, 1995; Laiti et al.,
1998; Eng et al., 2000), and, given sufficient time, to gibbsite (a-Al(OH)3), the
thermodynamically stable, fully hydrated alumina phase. Thus the surfaces of
aluminum oxides used in aqueous sorption experiments are not likely to have the
same structures or compositions as the anhydrous starting material. Even though
there have been hundreds of studies of the interaction of water with clean metal
and metal-oxide surfaces over the past 30 years using a variety of surface science
methods (e.g., Thiel and Madey 1987; Henderson 2002), there is little fundamental
understanding of how liquid water reacts with mineral surfaces under relevant
environmental conditions except in the simplest cases (Brown, 2001).
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Maureen and Bruce Garrett, also a quantum chemist at PNNL, organised a small,
focused workshop in 1992 in Seattle, to bring together three experimentalists
(George Parks and me from Stanford University and Jim Morgan from Caltech) and
about 20 theoreticians to discuss environmentally relevant chemical reactions on
mineral surfaces. I recall that this conversation took place on the final morning
of the workshop during the mid-morning coffee break. Over the previous two
days of the workshop, interactions between the three experimentalists and the
theoreticians were analogous to ships passing in the night, figuratively speaking.
The experimentalists did not understand the theoreticians and vice versa. One of
the talks by a well-known theoretician focused on a simulation of liquid argon
between two parallel plates of mica. I remember making the comment following
this talk that I could not remember ever seeing a pool of liquid argon in the environment. Another talk by a theoretician reported the largest MD simulation up
to that time of Na+ and Cl ions in a box of over 1,000 water molecules. The final
result of these heroic calculations was that these ions diffused to the boundaries
of the box, suggesting that their interaction with water was hydrophobic. Looking
back, it is now clear that the ab initio interatomic potentials available at that
time for that MD simulation did not allow accurate modelling of the chemical
interactions between water molecules and these ions. I left that workshop with
a renewed commitment to understand how water reacts with clean metal-oxide
surfaces. Ping Liu, my graduate student, and Tom Kendelewicz, a senior research
associate in my group, refurbished an old set of coupled UHV chambers at SSRL
in preparation for the experiments, and Tom and I wrote a proposal for beam time
on SSRL beam line 10-1. I also contacted Lynn Boatner of the Solid State Division
at Oak Ridge National Lab about single crystal samples of various metal oxides
he had grown years earlier as part of another project. What a lucky call that was.
Lynn had stored many MgO and CaO crystals he had synthesised years earlier in
mineral oil and sent the entire batch to me for use in our experiments. Our SSRL
proposal was approved as was an NSF Earth Sciences proposal, and we began a
five-year study that resulted in some new insights about this important reaction.
Because MgO has perfect (100) cleavage, we prepared (100) surfaces of
MgO single crystals by cleaving them in a UHV preparation chamber (base pressure of 10 10 Torr) just prior to the beginning of each XPS experiment. Each MgO
sample with a fresh (100) cleavage surface was then transferred to the analysis
chamber through a gate valve, and characterised by XPS survey scans, which
showed negligible adventitious carbon and no other impurities. The advantage
of using a synchrotron X-ray source rather than a laboratory XPS instrument,
which typically uses an Al or Mg sealed-tube X-ray source, is that the synchrotron X-ray source can be tuned to an energy that corresponds to the minimum
of the universal curve (a plot of the kinetic energy vs. escape depth of the excited
photoelectron from the sample surface; e.g., Lindau and Spicer, 1980). The result is
that higher surface sensitivity can be achieved with the synchrotron X-ray source.
Following collection of O1s XPS spectra for the clean surface, the sample was
transferred back to the preparation chamber and exposed for three minutes to
~10 8 Torr of ultrahigh purity water vapour using a precision leak valve. Following
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evacuation of the preparation chamber, the water-exposed sample was transferred back to the analysis chamber, and O1s, O2p, and valence band spectra
were taken. The same surface was exposed sequentially to increasing p(H 2O)
levels following this procedure, including, as a final step, immersion in liquid
water in a container in a glove box connected to the UHV preparation chamber.
One of the problems we encountered early in these experiments was
charging of the MgO surfaces during collection of the XPS spectra due to the
fact that MgO is a large band gap (~8 eV) insulator, resulting in a positive charge
build-up on the surface as photoelectrons are emitted. I remember calling up my
friend Vic Henrich, a surface physicist at Yale University and an expert on surface
science studies of metal oxides (Henrich and Cox, 1994), and asking him how to
solve this problem. It turned out that Vics XPS work had focused on metal oxides
that were small band gap semiconductors, and thus do not build up a positive
charge, so he was not able to help solve our problem. After more consultation
with others, we finally decided to employ a low energy electron flood gun to
compensate for the positive charging of the MgO(100) surface. We initially chose
an energy of 6eV for the flood-gun electrons, and later discovered, after discussions with Tom Orlando (Georgia Institute of Technology) that the secondary
electrons produced in the surface region of the MgO sample at this flood gun
energy were sufficiently energetic to cause some dissociation of water molecules
(Orlando et al., 1999). We had to repeat our XPS experiments up to that point at
a flood-gun energy of 4eV to avoid this problem. Following this learning experience, we obtained high quality O1s XPS spectra of the MgO(100) surface after
water vapour exposure.
Some of our O1s XPS results for the MgO(100)-water interaction are shown
in Figure 9.1a. Notice the growth of a shoulder on the low kinetic energy (high
binding energy) side of the O1s photopeak with increasing water content. This
shoulder represents hydroxyl groups from the heterolytic dissociation of water
on the MgO surface. Also notice that this shoulder does not increase much in
integrated intensity up to 10 4 Torr p(H 2O). However, slightly above this pressure
(3 10 4 Torr), this shoulder increased rapidly in integrated intensity. We interpreted this as evidence for dissociation of water initially on defect sites (step edges
and corners and oxygen vacancies) on the MgO(100) surface at p(H 2O)10 4
Torr, followed by dissociation of water molecules on terrace sites above 10 4
Torr p(H 2O). We also carried out a thermodynamic analysis of the interaction of
water with MgO(100) using the reaction of MgO with water to produce brucite
(Liu et al., 1998a):
(9.1)
o
which has a Gibbs free energy of reaction of -35.5 kJmol at 298.15 K, which was
the temperature of our sample in the XPS measurements, and assuming p(H2O) =
1 bar. The experimental pressure differs from that of the standard state (p(H 2O)0),
so we had to correct the pressure difference as follows:
p1
DGr = DG0r + p0 V dp
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(9.2)
(9.3)
The second term in equation (9.3) ((p1 1) DVs) is very small and can be neglected.
To find the lowest pressure at which the reaction occurs, we let DGr 0. Then
p(H2O),p1 exp {DG0r /(RT)} 5.9 10 7 bar = 4.5 10 4 Torr (at 298.15K) (9.4)
This calculated p(H 2O) is close to the value (3 10 4 Torr) at which we observed
a major increase in the integrated intensity of the high binding energy shoulder.
A similar calculation for hydroxylation of the CaO(100) surface gives p(H 2O) =
2.2 10 9 Torr at 298.15 oK, based on DGr = -65.9 kJ mol1 for the reaction CaO
+ H 2O = Ca(OH)2. This value is also close to the p(H 2O) at which CaO(100)
hydroxylates completely. When we were beginning the CaO(100) water dosing
experiments using the same procedure described above for MgO(100), we placed
a fresh single crystal of CaO in our preparation chamber, pumped down to a base
pressure of ~10 10 Torr, and cleaved the CaO crystal. We looked at the time and
listened to the grumbling in our stomachs and decided to take a break for dinner
before starting the water dosing experiment on CaO(100). Following a leisurely
dinner at a local Mexican restaurant, we returned to the beam line at SSRL and
resumed the experiment. We transferred the clean CaO(100) sample into the
analysis chamber and measured the O1s XPS. To our surprise, the CaO (100)
surface was fully hydroxylated even though we had not exposed it to water from
our precision water doser. Our conclusion was that there were sufficient water
molecules attached to the inner surface of the preparation chamber, even at a base
pressure of ~10 10 Torr, that the CaO (100) surface hydroxylated (Liu et al., 1998c).
Again, the observed p(H 2O) value was within about one order of magnitude to
that predicted from equilibrium thermodynamics.
These results led us to hypothesise that water dissociatively chemisorbs
only at defect sites at very low p(H 2O) values, and that extensive hydroxylation
of terrace sites does not occur until a threshold p(H 2O), above which complete
hydroxylation of the surface takes place. The nature of these interactions is
shown schematically in the structural drawings in Figure 9.1a, which depict two
MgO(100) surfaces, one showing terrace sites, a step defect, and a vacancy defect
prior to exposure to water vapour, and the second showing hydroxyl groups on
the surface following exposure at p(H 2O)>3 10 4 Torr. At p(H 2O) values less
than this threshold p(H 2O), water chemisorbs dissociatively on these types of
defects, which are thought to be more energetic and reactive sites than terrace
sites on MgO(100) based on DFT calculations (Langel and Parinello, 1994; Scamehorn et al., 1994). The associated low kinetic energy shoulder in the XPS spectra
does not grow appreciably in intensity as p(H 2O) is increased from 10 9 to 10 4,
even after prolonged exposures (>6hr) of the surface to water vapour (1.8
104 Langmuirs, L; where 1 L corresponds to 10 6 Torr sec). However, when the
defect density of these surfaces was increased (from 3% to 35%; Fig. 9.1b), this
low kinetic energy shoulder increased significantly in intensity at p(H 2O)<10 3
Torr (Liu et al., 1998b). In addition, when p(H 2O) was raised to 3 10 4 Torr with
an exposure time of 3 minutes (corresponding to an exposure of 5.4 104 L), the
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Figure 9.1
(a) Crystal structure of the clean MgO (100) surface (bottom), the same structure upon first exposure to water molecules
showing the hydroxyl groups at the step edges and oxygen vacancy site (middle) and the crystal structure of brucite
(Mg(OH)2), which is the fully hydroxylated product of the MgO-water reaction and selected O1s XPS spectra of MgO(100)
at different water vapour exposures from 2.3 10 8 Torr to 10 3 Torr and finally full immersion in water (after Liu et al.,
1998a). (b) O1s XPS of clean MgO(100) and MgO (100) surfaces with various levels of defects (3%, 18%, 35%) after exposure
to p(H2O) of 2-3 10 8 Torr (after Liu et al., 1998b).
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et al., 2000) are in substantial agreement with our photoemission results (Liu
et al., 1998d). Similar XPS experiments on -Fe2O3(0001) indicate a threshold
p(H 2O) of 10 4 Torr, above which water dissociates on terrace sites (Liu et al.,
1998d). These observations raise the question as to why the threshold p(H 2O)
values of isostructural corundum and haematite differ by about five orders of
magnitude (see Section 10).
Based on these and other ex situ metal oxide-water experiments, we found
that the threshold p(H2O) varies for the different metal-oxide surfaces examined
as follows: MgO(100) = 1 10 4 Torr, CaO(100) 5 10 10 Torr, a-Al2O3(0001) = 1
Torr, a-Fe2O3(0001) = 1 10 4 Torr, Fe3O4(111) = 1 10 3 Torr, and TiO2(110)
0.6 Torr. What is clearly demonstrated by these and other similar studies is that
most freshly exposed metal-oxide surfaces react rapidly with liquid water or water
vapour in the atmosphere and become fully hydroxylated. An important question
addressed below is how surface hydroxylation affects the structure/reactivity of
metal-oxide surfaces.
Figure 9.2
Oxygen 1s XPS of a-Al2O3 (0001) (a) and a-Fe2O3 (0001) (b) as a function of
p(H2O) (in Torr; after Liu et al., 1998d).
The metal oxide-water XPS studies described above were ex situ experiments in the sense that the metal-oxide surfaces were exposed to water in a
preparation chamber, which was pumped down to a base pressure of <10 10 Torr
following each water exposure, then the water-reacted sample was transferred
to the analysis chamber for XPS measurements. In 1997, David Shuh (LBNL)
and I (GB) led an effort to develop the scientific case for a new soft X-ray/VUV
Molecular Environmental Science beam line at the ALS (Brown, 1997). We were
successful, and the DOE BES Chemical Sciences Division provided the funding
to build two branch lines on an elliptical undulator source (BL 11.0.2) one
devoted to ambient-pressure XPS and one devoted to scanning transmission
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X-ray microscopy. One of the justifications for the ambient-pressure XPS branch
line was that it would allow us to study the reaction of water and other liquids
and gases with metal oxides under in situ conditions. As soon as the beam line
was completed, my group and the groups of Anders Nilsson (SLAC National
Accelerator Laboratory), Hendrik Bluhm and Miquel Salmeron (both LBNL) began
a collaborative effort, with our first projects designed to re-investigate the interaction of water vapour at RHs up to ~34% with the (0001) surface of -Fe2O3
(Yamamoto et al., 2010).
Figure 9.3a shows these O1s XPS results under in situ conditions (i.e. the
spectra were taken while the surface was in contact with water vapour at RHs up
to ~35%). The surface was prepared by cutting and polishing a single crystal of
a natural haematite (specularite) from Bahia, Brazil. After polishing, the sample
was acid-etched with 0.2 M HNO3 solution and then rinsed with Milli-Q water.
After introduction into the UHV chamber, the sample was cleaned by several
cycles of annealing at 723-773oK in 1 10 5 % O2. The cleaned surface displayed
a sharp (1 1) low energy electron diffraction pattern, indicative of a high degree
of surface order. We found that hydroxylation of the haematite surface begins at
the very low RH of 10 7 % (Fig. 9.3b) and precedes the adsorption of molecular
water. The coverage of OH increases with an increase in RH, and increases more
rapidly after the onset of water adsorption, which was attributed to a cooperative effect among adsorbed water molecules. This water-catalysed dissociation
is explained by the stabilisation of the dissociated final state due to the strong
hydrogen bond between H 2O and OH, which lowers the kinetic barrier for water
dissociation (Andersson et al., 2008). At high RHs the surface is covered with 1
monolayer (ML) of OH species and molecular water adsorption increases rapidly
after OH coverage reaches one ML. This maximum coverage of OH is consistent
with previous studies where 1 ML of hydroxyl species was observed on haematite powders exposed to ambient pressure water vapour (e.g., McCafferty and
Zettlemoyer, 1971) and on -Fe2O3(0001) single-crystal surfaces exposed to air
or immersed in bulk water (Junta-Rosso and Hochella, 1996).
We have extended our in situ XPS studies to MgO(100) (Newberg et al.,
2011a,b) and Fe3O4(100) (Kendelewicz et al., 2012), with results similar to those
described above. Our ex situ and in situ XPS and ab initio thermodynamic results
discussed above, indicate that water chemisorbs dissociatively only at defect sites
on these surfaces when at very low coverages of less than ~0.2 ML. At a certain
critical coverage where water molecules and surface OH groups interact through
H-bonding (a cooperative effect), corresponding to a threshold p(H 2O), water
chemisorbs dissociatively on terrace sites, resulting in the build-up of ~1 ML of
OH on the metal-oxide surface and more than 1 ML of water molecules adsorbed
on the surface hydroxyl layer.
GEOCHEMICAL PERSPECTIVES | G O R D O N E . B R O W N J R . G E O R G E S C A L A S
573
Figure 9.3
574
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GEOCHEMICAL PERSPECTIVES | G O R D O N E . B R O W N J R . G E O R G E S C A L A S
575
simple termination of the bulk structure along the (1-102) plane has only oneand three-coordinated oxygens. These observed differences in structure of the
hydrated (0001) and (1-102) alumina surfaces help explain the greater reactivity
of the latter to certain metal ions in solution, as will be discussed in Section 11
on Pb2+ sorption products.
Figure 10.1
576
GEOCHEMICAL PERSPECTIVES | V O L U M E 1 , N U M B E R 4 A N D 5
phase (e.g., O2 or H 2O) in terms of the Gibbs free energy of the solid surface as a
function of T and P and the number and chemical potentials of the atoms in the
system. g(T,P) is normalised to energy per unit area which can be expressed as:
(10.1)
where A is the surface area of the solid (typically modelled as a slab of the structure with two equivalent surfaces), G is the DFT-computed Gibbs free energy of
the simulation cell at pressure p and temperature T, and Ni and mi are the number
and chemical potential, respectively, of the different atoms in the three-dimensional supercell (hydrogens, metal atoms, and oxygens) at equilibrium with the
gas phase reservoir. Reuter and Scheffler related the DFT total energy, evaluated
for a certain volume of the unit cell, to the Gibbs free energy of the system by
including a vibrational energy term that is based on the phonon density of states,
and take into account the difference in vibrational modes of the metal-oxide
surface and the bulk-metal oxide. The gas-phase chemical potentials are taken
from the NIST-JANAF Tables.
The first ab initio thermodynamic approach to environmental interfaces
combined with experimental determination of the structure of a hydrated metaloxide surface was carried on the hydrated a-Fe2O3(0001) surface (Trainor et al.,
2004). CTR diffraction measures the diffuse scattering between Bragg diffraction
maxima (Fig. 10.2), which is sensitive to surface structure. Various surface structural models are fit to this diffuse scattering, and a best-fit model is chosen (see
Robinson and Tweet, 1992 and Renaud, 1998 for a description of this method).
Our CTR analysis revealed that it differs significantly from the three possible
terminations of the bulk haematite structure (Fig. 10.3) and consists of roughly
equal numbers of hydroxyl groups coordinated by one, two, and three VI Fe(III),
as shown in Figure 10.4.
Our collaborator Anne Chaka carried out DFT-ab initio thermodynamic
calculations on haematite(0001) in the presence of water vapour, and the results
of Annes calculations of surface energies (in meV2) for the different structures
of haematite(0001) in equilibrium with water vapour are shown in Figure 10.5.
The (HO)3-Fe-O3-R termination represents the lowest energy surface in equilibrium with water vapour under the high p(H 2O)-low pO2 conditions of our
CTR experiments. The (HO)3-Fe-Fe-R (0001) termination is stable over a lower
range of pO2. (HO)3-Fe-O3 represents the three top layers of the (0001) surface
of hydrated haematite, with each outermost OH group bonded to a single Fe
atom in octahedral coordination and the second-layer Fe atom bonded to three
oxygens in the next layer down. (OH)3-Fe-Fe represents the three top layers of
the (0001) surface, with each hydroxyl in the outermost layer bonded to two Fe
atoms in the second and third layers, each of which is bonded to three oxygens
in the fourth layer.
GEOCHEMICAL PERSPECTIVES | G O R D O N E . B R O W N J R . G E O R G E S C A L A S
577
Figure 10.2
578
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Although the CTR best-fit results indicate that both O3-Fe-O3-R (40%)
and O3-Fe-Fe-R (60%) terminations are present (the protons are omitted because
they cannot be detected by X-ray scattering), we speculated that the presence of
the O3-Fe-Fe-R termination observed in the CTR fitting may be due to sample
preparation conditions and that the O3-Fe-O3-R termination is likely the most
stable one under the high p(H 2O)-low pO2 conditions of the CTR experiments.
Figure 10.3
(left panel) Atomic-layer sequence of the bulk unit cell of the -Fe2O3 structure
along the [0001] direction. (right panel) Surface layer models of the three
possible clean terminations: (top) the oxygen-terminated surface, where
O3-Fe-Fe represents the three outermost layers of the (0001) surface; (middle)
the double iron terminated (0001) surface; (bottom) the single iron terminated
(0001) surface (from Trainor et al., 2004, with permission from Elsevier).
GEOCHEMICAL PERSPECTIVES | G O R D O N E . B R O W N J R . G E O R G E S C A L A S
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Figure 10.4
Atomic-layer sequence along the [0001] direction showing the relaxed models
of the two hydrated -Fe2O3 (0001) surfaces that best fit the CTR data (Fig.
10.2). The larger (red) and smaller (blue) spheres represent O and Fe atoms,
respectively. A layer of water molecules is shown above each surface. The
top panel shows the O3 -Fe-Fe-R surface, and the bottom panel shows the
O3-Fe-O3-R surface (from Trainor et al., 2004, with permission from Elsevier).
Figure 10.5
Surface energy vs. oxygen chemical potential for different haematite (0001)
terminations. The vertical dashed line on the left is the binding energy of
oxygen in the bulk structure, and the vertical dashed line on the right corresponds to the energy of an O2 condensate (after Trainor et al., 2004).
580
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GEOCHEMICAL PERSPECTIVES | G O R D O N E . B R O W N J R . G E O R G E S C A L A S
581
Figure 10.6
Figure 10.7
582
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in equilibrium with water based on solvaton theory and crystal chemistry considerations (specifically Paulings electrostatic valence rule) and Table 10.1 lists a
number of minerals, their bulk dielectric constants, ek, and their experimentally
measured and calculated pH PZC values.
Table 10.1
Phase
ek
pHpzc (exp)
pHpzc (calc)
-SiO2
SiO2 (am)
4.58
2.9
2.91
3.81
25.0
3.5
8.5
3.9
8.47
-Fe2O3
-FeOOH
Fe3O4
11.7
9.0-9.7
9.4
20 000
10.4
6.6
9.1
7.1
9.37
-Al2(OH)3
8.4
10.0
9.84
-TiO2
120.9
5.8
5.2
b-MnO2
ZrO2
UO2
MgO
CaO
NiO
CuO
Kaolinite
Muscovite
Microcline
Lo Albite
Hi Albite
10 000
4.6-7.3
4.8
22.0
24.0
9.6
11.95
11.9
18.1
11.8
7.6
5.5
6.95
6.95
12.4
9.85-11.3
9.5
4.5
6.8
2.0
-Al2O3
7.9
9.2
12.24
12.3
11.8
8.6
4.66
6.6
6.1
5.2
2.8
Among the common mineral sorbents, manganese and iron oxides play a
prominent role in sorbing first-row transition metal ions and heavy metal ions
from natural waters (Usui, 1979; Li, 1982; Stumm and Morgan, 1996). They are
particularly important sorbents in the deep ocean, where they occur as ferromanganese oxide nodules, and on seamounts where they occur as crusts. According
to Li (1982), Mg2+, Ba 2+, Ni2+, Cu2+, Zn2+, and Cd2+, all with low-to-moderate first
hydrolysis constants, are preferentially associated with Mn in deep-sea nodules,
whereas Co2+ and Pb2+ are often associated with Fe in these nodules.
Takematsu (1979) and Li et al. (1984) determined the relative affinities of
divalent cations in seawater for goethite, vernadite, and aluminosilicate phases
(clays) and found that vernadite has the highest Kd and that pelagic clay minerals
GEOCHEMICAL PERSPECTIVES | G O R D O N E . B R O W N J R . G E O R G E S C A L A S
583
have the lowest. Usui (1979) suggested that the affinity of Co2+ and Pb2+ for
ferromanganese crusts may be due to the ready oxidation of these cations on the
surfaces of manganese oxides, which would probably preclude their incorporation
into the tunnel structure of todorokite, but would enhance their sorption onto
HFOs. The higher affinity of many cations for vernadite and goethite vs. clay
minerals is attributed to the smaller intrinsic acidity constants of vernadite and
goethite relative to clay minerals (for the reaction -X-OH = -X-O + H+) and to
the higher dielectric constants of vernadite (~32) (Murray, 1975) and goethite vs.
clay minerals (4.5 to 8) (Keller, 1966), resulting in a smaller DGsolvation for Mn and
Fe oxides than for clay minerals (Takematsu, 1979; Li et al., 1984).
Perspectives on Hydrated Mineral Surfaces and Macroscopic Studies of Sorption Processes on These Surfaces The macroscopic uptake studies of aqueous
cations and anions on mineral surfaces, pioneered by Werner Stumm and his students,
have shown significant differences in the reactivity of different mineral surfaces with
respect to a given ion. They have also shown that different cations and anions have
significantly different reactivities to a given mineral surface. The synchrotron-based
surface X-ray scattering studies of hydrated mineral surfaces discussed above have
shown that the surfaces of nominally anhydrous metal oxides such as haematite and
corundum are not simple terminations of the bulk structure, but undergo significant
relaxations or reconstructions when in contact with water. The results of these studies,
coupled with DFT-ab intio thermodynamic studies of the same surfaces, also provide
molecular-level explanations of the differences in chemical reactivity of metal-oxide
surfaces to aqueous cations and anions. The future challenge is to extend these
types of combined macroscopic uptake and x-ray structural studies of both sorption
complexes and hydrated surfaces to other important mineral surfaces, including the
rock-forming silicates.
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11.
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585
Name
Microcline
Anglesite
Cerrusite
Hydrocerrusite
Galena
Laurionite
Leadhillite
Litharge
Massicot
Plumbojarosite
Plumbogummite
Pyromorphite
Vanadinite
586
Formula
(K,Pb)AlSi3O8
PbSO4
PbCO3
Pb3(CO3)2(OH)2
PbS
PbOHCl
Pb(SO4)(CO3)2(OH)2
PbO (yellow)
PbO (red)
Pb[Fe3(SO4)2(OH)6]2
PbAl3(PO4)2(OH)5H2O
Pb3(PO4)2Cl
Pb3(VO4)3Cl
Relative solubility
Relatively Soluble
Soluble
Soluble
Insoluble at high pH
Insoluble
Soluble
Soluble
Soluble
Soluble
Soluble
Highly Insoluble
Highly Insoluble
Insoluble
GEOCHEMICAL PERSPECTIVES | V O L U M E 1 , N U M B E R 4 A N D 5
GEOCHEMICAL PERSPECTIVES | G O R D O N E . B R O W N J R . G E O R G E S C A L A S
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588
GEOCHEMICAL PERSPECTIVES | V O L U M E 1 , N U M B E R 4 A N D 5
Figure 11.1
(11.2)
where sOH is the bond valence of the O-H bond (in valence units) and ROH is the
O-H distance (in ). We used this relationship to predict the bond valences of
O-H bonds to surface oxo groups on the -Al2O3 and -Fe2O3(0001) and (1-102)
surfaces. Note that O-H bonds characteristic of hydroxyl groups can vary from
0.95 to 1.03 , resulting in O-H bond valences 0.88 to 0.68 v.u., respectively,
and that OH bonds associated with the acceptor bonds of water molecules
can vary from ~1.65 to ~2.50 , with corresponding O H bond valences
of ~0.25 to ~0.13 v.u., respectively (Bargar et al., 1997a). From our CTR studies
of hydrated surfaces, we now know that the -Al 2O3(1-102) surface (Trainor
et al., 2002c) has roughly equal numbers of oxygens coordinated by one, two,
and three VI Al3+ ions and that the -Al2O3 (0001) surface (Eng et al., 2000) has
oxygens coordinated to two VI Al3+ ions, dominantly. The EXAFS spectroscopy
results of Bargar et al. (1997a) indicate that Pb2+(aq) is coordinated by 3 oxygens,
GEOCHEMICAL PERSPECTIVES | G O R D O N E . B R O W N J R . G E O R G E S C A L A S
589
Figure 11.2
(a) Plots of bond valence of O-H bonds (sOH) vs. O-H distance (R in ) after
Bargar et al. (1997a) and Brown (1987). (b) Schematic illustration of Pb2+ and
Co2+ adsorption complexes on the -Al2O3(1-102) and (0001) surfaces, showing
bond-valence sums (Ss) to the surface oxo and hydroxo groups. The larger open
circles represent oxygens, the dark grey circles represent Al3+ ions, and the
small black circles represent protons. Pb2+ and Co2+ are represented by light
grey circles, which are labelled (after Bargar et al., 1997a).
Another important finding from the Bargar et al. (1996, 1997a) studies is that
no mixed-metal Pb2+-Al3+-hydroxide phase (e.g., the hydrotalcite structure type)
forms under the experimental conditions. This finding is in contrast with that
for Co2+, Ni2+, and Zn2+ sorption at the alumina-water interface, where relatively
590
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GEOCHEMICAL PERSPECTIVES | G O R D O N E . B R O W N J R . G E O R G E S C A L A S
591
Figure 11.3
(a) Pb LIII XANES spectra of Pb2+ adsorbed at the -Fe2O3 (0001) and (1-102)
and -Al2O3(1-102) surfaces in contact with water. (b) Background-subtracted,
k 3-weighted EXAFS spectra of Pb2+ sorbed on the -Fe2O3 (0001) and (1-102)
surfaces compared with Pb2+ adsorbed on an -Fe2O3 powdered sample and
Pb 4(OH) 4 4 aqueous complexes. (c) Fourier transforms of the EXAFS spectra
shown in the middle panel (after Bargar et al., 2004).
592
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GEOCHEMICAL PERSPECTIVES | G O R D O N E . B R O W N J R . G E O R G E S C A L A S
593
Figure 11.5
(a) As K-edge XANES recorded at 10 oK for As(III) sorbed onto biogenic magnetite nanoparticles as a function of As(III) surface coverage. (b) Backgroundsubtracted k 3-weighted EXAFS spectra of these samples. (c) Magnitude and
imaginary part of the Fourier transform (FT) of these EXAFS spectra. Note the
decrease of the second-neighbour contribution to the EXAFS with increasing
surface coverage. Experimental and calculated k 3-weighted EXAFS as well as
corresponding FT curves are displayed as dashed and solid lines, respectively
(from Morin et al., 2009, with permission from the American Chemical Society).
Figure 11.6
(a) EXAFS-derived structural information for the As(III) surface complexes. (b)
Proposed structural model for the As(III) tridentate, hexanuclear, corner-sharing
complexes (3C) on the (111) surface of magnetite; As(III)O3 pyramids (black)
occupy vacant FeO 4 tetrahedral sites (dark grey). (c) Side view of 3 C As(III)
complex on a magnetite structural fragment. All As-Fe distances expected for
this complex are consistent with the EXAFS-derived distances (from Morin et
al., 2009, with permission from the American Chemical Society).
GEOCHEMICAL PERSPECTIVES | V O L U M E 1 , N U M B E R 4 A N D 5
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595
12.
Up to this point, we have focused mostly on XAFS studies of cation and oxoanion
sorption products at mineral-water interfaces and X-ray standing wave studies
of the EDL. Organics of various types are also important players in the EDL of
natural sorbents and are important sorbates in many environmental systems.
Here we summarise some of the results of our in situ (i.e. with liquid water present
at ambient T and P) ATR-FTIR spectroscopy studies of the sorption of simple
carboxylic acid molecules and natural organic matter (NOM) on model mineral
surfaces. ATR-FTIR is arguably the spectroscopic technique of choice for deriving
molecular-scale information on the attachment geometry of organic molecules at
solid-water interfaces (Hind et al., 2001) as demonstrated by Per Persson (Ume
University) and co-workers (e.g., Axe et al., 2006; Boily et al., 2000). In a series of
in situ ATR-FTIR studies, we examined the interaction of (1) oxalate (C2O42) with
-Al2O3 (corundum) and g-AlOOH (boehmite) (Yoon et al., 2004b; Johnson et al.,
2004b), (2) maleate (H 2C4O42) with -Al2O3 (Johnson et al., 2004a), (3) pyromellitate (C14H14O8) with -Al2O3 (Johnson et al., 2005a), (4) lactate (L-CH3CH(OH)
COO ) with -Fe 2O 3
(haematite) nanoparticles (Ha et al., 2008),
and (5) Suwannee River
fulvic acid and Pahokee
peat humic acid with
g -A lOOH ( Yoon et
al., 2004a, 2005b). We
also considered the
effect of the adsorption modes of maleate,
oxalate, and citrate
(C 3 H 5 O(COO) 3 3 ) on
the colloidal stability
of cor u ndu m-water
s u s p e n s ion s b a s e d
Figure 12.1 Schematic illustration of aqueous oxalate on electrokinetic and
complexes at a mineral-water interface showing shear-yield stress measboth aqueous complexes as well as inner- and
urements over a range
outer-sphere surface complexes.
of pH and organic acid
concentrations (Johnson et al., 2005b). Figure 12.1 is a cartoon of possible oxalate
complexes in solution and surface complexes at a mineral-water interface.
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In the Yoon et al. (2004b) study, at pH 5.1, at least four different oxalate
species were found at or near the boehmite-water interface for oxalate surface
coverages (ox) ranging from 0.25 to 16.44 mmol m2. At relatively low coverages
(ox <2.47 mmol m2), strongly adsorbed inner-sphere oxalate species (IR peaks
at 1286, 1418, 1700, and 1720 cm1) (Fig. 12.2) replace weakly adsorbed carbonate
species, and a small proportion of oxalate anions were found to be adsorbed in
an outer-sphere mode (IR peaks at 1314 and 1591 cm1). IR peaks indicative of
inner-sphere adsorbed oxalate were also observed for oxalate at the corundumwater interface at ox = 1.4mmolm2. With increasing oxalate concentration (ox >
2.47 mmol m2), the boehmite surface binding sites for inner-spherically adsorbed
oxalate become saturated, and excess oxalate ions were present dominantly as
aqueous species (IR peaks at 1309 and 1571 cm1).
T he coord i nat ion
geometry of inner-spherically adsorbed oxalate
species was also predicted
using quantum chemical
geometry optimisation and
IR vibrational frequency
calculations. Geometryopt im ised A l 8 O 12 and
Al14 O22 clusters with the
reactive surface Al site coordinated by three oxygens
were u sed a s model
substrates for corundum
and boehmite surfaces.
Among the models considered, calculated IR frequencies based on a bidentate
side-on structure with a
5-membered ring agree
best with the observed
frequencies for boehmite- Figure 12.2 ATR-FTIR spectra of oxalate in the presence
of boehmite (g-AlOOH) powder at different
oxalate-water samples at
oxalate sur face loadings (in molm 2 )
ox = 0.25 to 16.44 mmol m2
indicated by the different coloured lines,
and pH 2.5 and 5.1 and for
showing spectral features due to aqueous
a corundum-oxalate- water
oxalate species and both inner- and outersphere oxalate complexes at the boehmitesample at ox = 1.4 mmol
water interface (after Yoon et al., 2004b).
m 2 and pH 5.1. Based on
these results, we suggested
that oxalate bonding on boehmite and corundum surfaces results in 5-coordinated rather than 4- or 6-coordinated Al surface sites and that oxalate forms a
five-membered ring with a single surface oxo group (Fig. 12.3).
Geochemical Perspectives | G o r d o n E . B r o w n j r . G e o r g e s C a l a s
597
Figure 12.3
598
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Figure 12.4
Geochemical Perspectives | G o r d o n E . B r o w n j r . G e o r g e s C a l a s
599
(Fig. 12.7). At sub-monolayer surface coverages (SRFA = 1.20 and 2.20 mol m2),
several new peaks and enhancements of the intensities of a number of existing
peaks were observed. The latter spectral changes arise from several nonorganic
extrinsic species (i.e. adsorbed carbonate and water, for alkaline solution conditions), partially protonated SRFA carboxyl functional groups (near neutral pH
conditions), and small quantities of inner-spherically adsorbed SRFA carboxyl
groups and/or Al(III)-SRFA complexes (for acidic conditions). The spectra of
PPHA adsorbed at boehmite-water interfaces also showed changes generally
consistent with our observations for SRFA sorbed on boehmite.
Figure 12.5
Figure 12.6
600
GEOCHEMICAL PERSPECTIVES | V O L U M E 1 , N U M B E R 4 A N D 5
Figure 12.7
(a) In situ ATR-FTIR spectra of humic acid in aqueous solution. (b) In situ ATRFTIR spectra of humic acid in the presence of boehmite powder (after Yoon
et al., 2005b).
These results confirm that SRFA and PPHA are predominantly adsorbed at
the boehmite-water interface in an outer-sphere fashion, with minor inner-sphere
adsorption complexes being formed only under quite acidic conditions. They also
suggest that the positively charged boehmite-water interface stabilises SRFA and
PPHA carboxyl functional groups against protonation at lower pH. Measurements of the concentration of dissolved Al3+ ions in the absence and presence of
SRFA showed that the boehmite dissolution process is clearly inhibited by the
adsorption of SRFA, which is consistent with previous observations that outerspherically adsorbed organic anions inhibit Al-(oxyhydr)oxide dissolution.
Werner Stumm and his collaborators developed simple models for the interaction of organic molecules and other ligands with mineral surfaces that have
proven very useful in understanding the dissolution kinetics of metal oxides
(Furrer and Stumm, 1986; Zinder et al., 1986; Wieland et al., 1988; Stumm and
Wollast, 1990; Biber et al., 1994). For example, they found that organic molecules,
such as oxalate and salicylate, which form five- and six-membered rings with
surface functional groups, are efficient in enhancing dissolution rates (Furrer
and Stumm, 1986; Fig. 12.8). The results of these classic studies are consistent
with some of the more modern ATR-FTIR studies discussed above and serve as
an excellent illustration of the importance of coordination chemistry in understanding chemical weathering or its inhibition in certain cases where organic
molecules form outer-sphere surface complexes (e.g., maleate or SRFA). Chapter
5 in Stumm (1992) and Sections 13.3 and 13.4 in Stumm and Morgan (1996)
provide excellent summaries of Stumms ideas on the kinetics and mechanisms
of chemical weathering of minerals.
GEOCHEMICAL PERSPECTIVES | G O R D O N E . B R O W N J R . G E O R G E S C A L A S
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Figure 12.8
602
GEOCHEMICAL PERSPECTIVES | V O L U M E 1 , N U M B E R 4 A N D 5
13.
The next step in our increasingly complex model of the mineral-water interface
addresses the question of how organic matter or microbial biofilm coatings impact
the reactions of mineral surfaces with trace levels of heavy metals. Close links
exist among biological activity, the organic matter generated by this activity, and
the mineralogy and geochemistry of the elements involved in low-temperature
processes. Biological activity has an enormous impact on mineral-water interfacial processes. For example, bacteria accelerate Fe oxidation reactions in iron
sulphides and arsenides (e.g., Vaughan and Lloyd, 2011), whereas other microorganisms favour the reduction of Fe(III) in Fe-(oxyhydr)oxides (e.g., Hansel et
al., 2003). Microbial organisms most commonly occur in consortia known as
biofilms, in which the bacteria are embedded in a hydrated matrix of extracellular
polymeric substance (EPS) that adheres to surfaces (Geesey et al., 1978; Allison
and Sutherland, 1987; Marshall, 1992; Costerton et al., 1995). Such biofilms are
widespread in soils and form microenvironments in which aqueous chemical
conditions differ from those of the host groundwater. The low isoelectric points
of bacterial surfaces (Busch and Stumm, 1968) and the abundance of anionic
functional groups on both bacterial surfaces and in the EPS produced by bacteria
(Beveridge and Murray, 1980; Sutherland, 1985; Beveridge, 1989) create a variety
of binding sites for metal ions. As a result, bacterial and polymer surfaces have
a high affinity for metal ions even at low pH, and uptake is enhanced at neutral
pH (Rudd et al., 1984; Ferris et al., 1989; Kellems and Lion, 1989; Geesey and Jang,
1990; Southam et al., 1995). Metal sorption data indicate a distribution of metal
binding sites consisting mostly of carboxyl and phosphoryl functional groups
(Beveridge and Murray, 1980; Fein et al., 1997; Daughney and Fein, 1998). Initial
binding of metal ions may occur on reactive sites within the bacterial cell wall
(Beveridge and Murray, 1980; Collins and Stozky, 1992) where the adsorption of
metals is rapid and reversible (Kellems and Lion, 1989; Collins and Stozky, 1992).
The bound metals may then act as nucleation sites in the formation of silicates,
carbonates, phosphates, sulphides, and organo-metallic complexes containing the
metal ion (Beveridge et al., 1983; Ferris et al., 1987, 1989; Thompson and Ferris,
1990; Urrutia and Beveridge, 1993, 1995; Fortin and Beveridge, 1997; SchultzeLam et al., 1992; Templeton et al., 2003a,b). The initial association of silicate,
sulphate, phosphate, and carbonate anions might be as outer-sphere complexes
bridged through bound multivalent metal cations (Schultze-Lam et al., 1992;
Urrutia and Beveridge, 1993) or by electrostatic attraction to the small number of
available positively charged functional groups (e.g., amine groups) (Beveridge and
Murray, 1980). Warren and Ferris (1998) showed that a continuum exists between
metal sorption and precipitation reactions on bacterial surfaces and that these
processes can be described by the same surface complexation/surface precipitation theory applied to sorption reactions at mineral surfaces. Such processes have
major implications for ground water quality, as the migration and toxicity of heavy
GEOCHEMICAL PERSPECTIVES | G O R D O N E . B R O W N J R . G E O R G E S C A L A S
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angles with increasing [Pb] indicates that Pb2+ is also binding to sites in the B.
cepacia biofilm coating. At all Pb concentrations studied, the FY-data indicate that
Pb2+ binds primarily to functional groups in the biofilm on the -Al2O3 (0001)
sample. The results of this study show that Pb2+ binds initially to reactive sites on
-Al2O3(1-102) and -Fe2O3(0001) even with a biofilm coating that covers essentially the entire mineral surface, as confirmed by other microscopic studies. The
order of reactivity of these biofilm-coated surfaces for Pb2+(aq) [-Fe2O3(0001)>Al2O3(1-102)>>-Al2O3(0001)] is the same as that observed in uptake and EXAFS
studies of Pb2+ sorption on biofilm-free alumina and haematite surfaces (Bargar
et al., 1996; 1997a,b,c; 2004), which is also consistent with the general findings of
Brydie et al. (2004, 2009).
Figure 13.1
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Figure 13.2
Measured (dashed) and modelled (light line) reflectivity (Log I 0/I1) profiles
and Pb L FY profiles (circles) with model fits (heavy line) for -Al2O3(0001),
-Al2O3 (1-102), and -Fe2O3 (0001) at (a) 10 6 M and (b) 10 5M [Pb] (after
Templeton et al., 2001).
Bacteria
Shewanella oneidensis MR-1
(Ha et al., 2010)
B. subtilis
(Fein et al., 1997)
B. licheniformis
(Daughney and Fein, 1998)
Types
of binding
sites
Site
concentrations
4.6
3.7
4.2
5.6
4.7
5.7
Carboxyl
Phosphoryl
Carboxyl
Carboxyl
1.4
606
Binding
affinities
(log Kapp)
102 mol/gdry
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3.4/8.75
7.00
Table 13.1
Mineral
shbstrates
Types of
binding sites
Site densities
(log n)1
-Al2O3 (0001)
M1
M2
6.25
5.05
-Al2O3 (1-102)
M1
M2
6.25
5.15
-Fe2O3 (0001)
M1
M2
5.85
5.1
Binding affinities
pHpsc
for Pb(II) (log
(point of
Kapp)1
zero charge)
5.35
4.12
2.7
6.0
5.22
3.55
6.65
6.43
3.5
Table 13.2
Figure 13.3
LP-X SW- F Y spec trum of Pb 2+ on a polyacr ylic acid (PA A) - coated (a)
-Al2O3(0001), (b) -Al2O3(1-102), and (c) -Fe2O3(0001) substrate, showing
in green the FY component from Pb in the PAA coating and in red the FY
component from Pb at the metal-oxide surface (after Yoon et al., 2004a).
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MR-1 biofilm coatings and metal-oxide substrates (Figs. 13.4 and 13.5), show that
the biofilms do not block all reactive sites on the alumina and haematite surfaces
and that sites on -Al 2O3(1-102) and -Fe2O3(0001) outcompete functional
groups in the biofilm (including the EPS exudate) at low Pb concentrations. These
results challenge the generalisation that NOM or biofilm coatings change the
adsorption characteristics of mineral surfaces (cf. Neihoff and Loeb, 1974), and
they are also inconsistent with the suggestion that NOM coatings block reactive
sites on mineral surfaces (cf. Davis, 1984).
Figure 13.4
(left) LP-XSW-FY spectra (open circles), fits (red line), and X-ray reflectivity data
blue circles and fits for Pb2+ at different lead concentrations on Shewanella
oneidensis MR-1 biofilm-coated -Fe2O3(0001) substrate. (right) Plot of log
Pb2+ concentrations in the biofilm coating relative to Pb2+ concentrations on
-Fe2O3(0001) as a function of log Pb(aq) and the % Pb2+ in the biofilm and on
-Fe2O3(0001) at different Pb concentrations.
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Figure 13.5
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609
610
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(e.g., Madden and Hochella, 2005). An example of this effect is seen in the rapid
increase in rate of Mn2+ oxidation by haematite nanoparticles at particle sizes
<10 nm (Fig. 14.1).
The reactivity of metal-oxide nanoparticles also can change with decreasing
particle size. This was shown to be true for TiO2 nanoparticles, which, relative to bulk TiO2, exhibit differences in photocatalytic reduction of cations like
Cu 2+ and Hg2+ in the presence and absence of surface adsorbers like alanine,
thiolactic acid, and ascorbic
acid due to different surface
structures (e.g., Rajh et al.,
1999). Shorter Ti-O bonds
and increasing disorder
around Ti with decreasing
size of the TiO2 nanoparticles suggest that the unique
surface chemistry exhibited
by nanoparticulate TiO2 is
related to the increasing
number of coordinatively
unsaturated surface Ti sites
with decreasing nanopar Figure 14.1 Rate of Mn 2+ oxidation vs. haematite ticle size (Chen et al., 1999).
nanoparticle size (from Hochella et al., The structure of amorphous
2012, with permission from Pan Stanford
TiO2 nanoparticles has been
Publishing).
modelled using a highly
distorted outer shell and a small, strained anatase-like core with undercoordinated or highly distorted Ti-oxygen polyhedra at the nanoparticle surface (Zhang
et al., 2008). Similarly, in nanoparticulate a-Fe2O3, surface Fe sites are also undercoordinated relative to Fe in the bulk structure; they are restructured to octahedral sites when the nanoparticles are reacted with enediol ligands (Chen et al.,
2002), which may explain an increased sorptive capacity for Zn2+(aq) relative to
larger-sized haematite particles found in our work on haematite nanoparticles
(Ha et al., 2009), although care must be taken to show that nanoparticulate surface
precipitates are not responsible for the increased sorption. For example, Morin
et al. (2009) found that precipitation of amorphous Fe-As-containing hydroxide
nanoparticles on nanocrystalline magnetite explained the larger-than-normal
uptake of As on magnetite (Section 11.2).
Another example is ZnS nanoparticles, which are often associated with
acid mine drainage environments and have been the focus of several structural
studies. A pioneering study of 3.4 nm ZnS nanoparticles (Gilbert et al., 2004)
found that structural coherence is lost over 2 nm and that the structure of the
nanoparticle is stiffer than that of bulk ZnS, based on a higher Einstein vibration
frequency in the nanoparticle. The surface region of the nanoparticle is highly
strained. In a similar study of ZnS nanoparticles in contact with aqueous solutions
containing various inorganic and organic ligands, stronger surface interactions
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with these ligands result in a thicker crystalline core and a thinner distorted outer
shell (Zhang et al., 2010). These and other studies strongly suggest that the surface
structures of certain types of nanoparticles are different than those of their larger
counterparts, so structural differences as well as the increased surface areas of
nanoparticles are factors in their enhanced chemical reactivities.
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Figure 14.2
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Figure 14.3
(left) X-ray scattering data for ferrihydrite with progressive aging at 175 C in
the presence of citric acid, showing the change in magnetic susceptibility, prior
to and after transformation to haematite. (upper right) Proposed structure
of 11-line ferrihydrite, showing the magnetic moments on the individual iron
atoms in a ferrihydrite unit cell. (lower right) Plots of the strain of the a and
c cell parameters of ferrihydrite as a function of particle size (after Michel
et al., 2010).
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coordinated Fe3+ (Eggleton and Fitzpatrick, 1988, 1990; Manceau et al., 1990). This
discussion, which has lasted 24 years, is another example of how the scientific
method should work (see another example in Section 8).
My student Cristina Cismasu recently completed her Ph.D. work on the
structure and reactivity of natural ferrihydrites, which included (1) detailed characterisation studies of natural ferrihydrites from an acid mine drainage system
associated with the second largest mercury mine in North America (located in
New Idria, central California) (Cismasu et al., 2011; Fig. 14.4), (2) structural studies
of the structure of synthetic ferrihydrites prepared in the presence of 0-40 mole%
Al (Cismasu et al., 2012) and 0-40 mole% Si (Cismasu et al., submitted-a), and
(3) macroscopic uptake and XAFS spectroscopic studies of natural and synthetic
ferrihydrites reacted with Zn 2+(aq) (Cismasu et al., submitted-b). As shown in
Figure 14.4, ferrihydrite in the New Idria acid mine drainage system is intimately
associated with NOM at the nm scale, with a variety of organic C functional
groups present as well as carbonate. The organic carbon comes in part from
breakdown of the extensive biofilms that reside on the surface of the water in
the acid mine drainage (AMD) pond located below the main mine waste dump,
where marcasite is undergoing oxidation and causing acidic pH values (2.5 to 3.5).
Figure 14.4
616
(left) STXM maps of Fe and C for two samples (NIFh1 and NIFh2) of ferrihydrite
collected at the New Idria acid mine drainage site. (right) Carbon K-edge
XANES spectra taken at the locations shown by the blue, red, and green circles
using a 30nm diameter X-ray beam, showing different carbon functional
groups including aromatic (~285 eV), phenolic (~286.5 eV), aliphatic (~287.5
eV), carboxylic (~289 eV), and carbonate (~290.5 eV) for the two samples; the
elemental maps were made by taking images below and above the C K- and
Fe LII-absorption edges, and subtracting the two images to produce maps with
enhanced sensitivity to C and Fe (after Cismasu et al., 2011).
GEOCHEMICAL PERSPECTIVES | V O L U M E 1 , N U M B E R 4 A N D 5
Similar STXM elemental maps were made for Al and Si and showed
many discrete submicron hotspots due to the formation of physically separate
phases of Al-hydroxide (most likely a gibbsite-like phase based on Al K-edge
XANES spectra) and SiO2 intergrown with the ferrihydrite at the submicron
scale (Cismasu et al., 2011). Cristina also collected high energy total X-ray scattering data at the APS on these natural ferrihydrite samples from New Idria and
Fourier transformed these very high Q data (~30 1) to produce high-resolution
PDFs (plotted as G(r) vs. R at right in Fig. 14.5). Also shown in Figure 14.5 are the
various Si-O, Fe-O, and Fe-Fe distances resulting from different SiO4 and FeOx
polyhedral linkages in ferrihydrite and a tabulation of these distances.
Figure 14.5
(right) PDFs for three natural ferrihydrite samples (NIFh1, NIFh2, NIFh3)
compared with that of synthetic ferrihydrite (Synth. fh). (left) Various Fe-O
and Fe-Fe distances within and between SiO 4 and FeOx polyhedra, responsible
for the features in the G(r) function; the table shows the various distances
observed for the different atom pairs (after Cismasu et al., 2011).
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2012). By comparing the amount of Al responsible for the NMR signal with the
total amount of Al present in these synthetic samples, Cristina was also able to
determine the amount of spectroscopically silent Al in the samples, which we
assigned to Al substituted for Fe3+ in the ferrihydrite structure. The lack of a
measureable NMR signal from these substituted Al ions is due to their proximity
to paramagnetic Fe3+ ions which broaden their NMR signals. This approach
showed that between 20 and 30 mole% Al substitutes for Fe3+ in these samples,
which compares well with other estimates of the extent of Al solid solution in
ferrihydrites (Cornell and Schwertmann, 2003).
Figure 14.6
(left) PDFs for synthetic ferrihydrites with different amounts of Si (in mol%);
coherent scattering domain size is determined by the distance (r) at which the
G(r) function no longer shows structure (i.e. where signal can no longer be
distinguished from noise). (middle) Models of the ferrihydrite structure for
various particle sizes after the Michel et al. (2010) structural model of ferrihydrite. (right) Plot of coherent scattering domain size vs. mole% Si for natural
ferrihydrite and synthetic Al- and Si-containing ferrihydrites.
The picture that emerges from Cristinas study of the effects of impurities on the structure of ferrihydrite is that natural ferrihydrites are often intimately associated with NOM and a variety of inorganic impurities, some at major
concentration levels (up to 20 mole% Si and Al), some substituting for Fe3+, such
as Al3+ and Cr3+, and some forming physically separate phases, such as Si4+
and some Al3+. This compositional complexity makes it difficult to sort out the
various factors controlling the reactivity of natural ferrihydrites because of the
number of different phases physically mixed at the nm-scale, all of which can
adsorb cations and oxoanions. Thus it is difficult to assess the reactivity of natural
ferrihydrites in these complex matrices and to determine the effects of impurity
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scattering, zeta potential measurements, and imaging. The results of our structural characterisation of the PVP-capped Ag-NPs (Fig. 14.8) show that there are
no discernable structural differences for the Ag-NPs of different sizes. These
results also show that there are no significant differences in the cell parameters
as a function of nanoparticle size, indicating that the nanoparticles have little if
any strain. As a result, we conclude that the surface regions of these nanoparticles
have a structure similar to that of the core regions.
Figure 14.8
(a) High energy total X-ray scattering data from Ag-NPs of different sizes.
(b) Fourier transforms of the X-ray scattering data, showing that there are
no significant structural differences for the different sized Ag-NPs. (c) Bestfit structural model of the Ag-NPs and cell parameters derived from the fits
(after Levard et al., 2011a).
Figure 14.7
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One of the key objectives of this project was to determine how organiccapped Ag-NPs transform under common environmental conditions, including
exposure to reduced sulphur. Clement exposed the Ag-NPs to different levels of
reduced S in the form of Na 2S(aq) and found that they quickly underwent dissolution followed by reprecipitation of acanthite (Ag2S), resulting in the formation of
Ag2S nanobridges between the Ag-NPs (Fig. 14.9).
Figure 14.9
TEM images of Ag-NPs (a) before and (b-f) after sulphidation at increasing
S/Ag ratios, showing the formation of acanthite (Ag2S) nanobridges between
the Ag-NPs (from Levard et al., 2011a, with permission from the American
Chemical Society).
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to be taken from Clements study is that ecotoxicology studies on the transformation products of manufactured nanoparticles in different environments are more
relevant than similar studies on the pristine nanoparticles.
Figure 14.10
622
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resulting from biological activity in aquatic systems (Buffle et al., 1998). In addition, unstructured gel-like NOM domains are associated with the biopolymers
or occur as isolated entities (Fig. 14.11, arrows).
Figure 14.11
TEM observation of river-borne particles from the Rio Negro. (a) Recomposed
image of a typical aggregate of biopolymers with gel-like organic domains
(arrows) and associated HFO. (b) Enlarged view of the dotted square region
in (a), showing an aggregate (top left) representative EDS spectra with Cu
originating from the grid. (centre) Fourier transforms of high-resolution
images, with interplanar distances indicated in nm (after Allard et al., 2004).
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626
GEOCHEMICAL PERSPECTIVES | V O L U M E 1 , N U M B E R 4 A N D 5
et al., 2003, 2007; Calas et al., 2008). In lateritic materials, the variation of defect
content of kaolinites matches the Fe concentration (Muller and Calas, 1989), and
the higher defect content of kaolinites in ferruginous nodules is related to the past
sorption of radionuclides on poorly ordered HFO, prior to their transformation to
haematite. Despite a limited penetration depth of ionising radiation (in the m
range for a-particles), radionuclide sorption on clays and associated oxides will
affect the crystal structure of clay platelets, preserving these ubiquitous witnesses
of radionuclide transfers in our environment. All of this work was possible due
to the experience of IMPMC on colour centres in minerals, allowing us to obtain
unique data on the behaviour of radionuclides in the environment.
Perspectives on Nanoparticles in the Environment Unraveling the structure
and properties of natural, incidental, and manufactured nanoparticles will continue
to be an active research area in Environmental Geochemistry and Mineralogy, particularly as novel approaches (e.g., Michel et al., 2010; Gilbert et al., 2010) are employed
and as we learn more about the ability of nanoparticles to facilitate the transport of
contaminants and pollutants and their health impacts on ecosystems as well as the
importance of colloids in the geochemical processes occurring during weathering
and erosion. Determining the structures of nanoparticles and nanoparticle surfaces
is one of the grand challenges of all branches of nanoscience because of the key role
that structure plays in determining properties such as reactivity. New developments
using single-particle scattering on X-ray free electron laser sources like the Linac
Coherent Light Source at the SLAC National Accelerator Laboratory are promising
(Bogan et al., 2008), although there is still much progress to be made by the next
generation of geochemists.
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and spent two weeks traveling through the Rocky Mountains and other areas
of Colorado, including Leadville. I was about 12 years old at that time and had
started collecting minerals a few years earlier. The mineralogy of Mississippi is
rather limited, with mostly sandstones, shales, mudstones, and other sedimentary
rocks and sediments, with no major economic mineral deposits other than sand,
gravel, clay, bentonite, chalk, and limestones, and with no underground mines
in the state. I was quite excited to see hard rocks and to have a chance to visit
some of the major economic mineral deposits of Colorado, including Leadville.
I recall collecting galena in the same waste dumps that I would work on years
later with my graduate student John Ostergren. John completed his undergraduate
degree in Geology at Carleton College in Minnesota then took a job with the
environmental consulting firm Walsh and Associates to assess the speciation of
lead in several tailings piles at Leadville. As John told me during his interview for
graduate school, he had reached a point in his work with Walsh and Associates
that was frustrating because he felt their approach was not fundamental enough.
I accepted John as a graduate student, and he joined my group in the fall
of 1993. John Bargar and I had begun doing XAFS spectroscopy studies of Pb2+
sorption products on haematite and corundum surfaces two years earlier, and
John Ostergren jumped at the chance to use XAFS spectroscopy and XPS, coupled
with selective chemical extractions (SCEs) and electron microprobe analysis, to
determine the speciation of lead and its potential bioavailability in two tailings
piles at Leadville one known as the Apache tailings that was rich in sulphides
and associated with acidic solutions, and one known as the Hamms tailings that
was carbonate buffered and associated with near-neutral pH solutions. Both tailings had about 8500 ppm total lead. The late Tracy Tingle (Research Associate,
Department of Geology, Stanford) joined us on this project, as did George Parks.
Figure 15.1 summarises some of the EXAFS results from Johns work (Ostergren et
al., 1999). The left panel shows a backscattered-electron image of a representative
thin section as well as the background-subtracted, k3-weighted EXAFS spectrum of the Apache tailings. Vanadinite, plumbojarosite, and plumboferrite were
identified by electron microprobe and in XRD analyses. The EXAFS spectrum
of the Apache tailings was fit using a linear combination of the EXAFS spectra
of these three crystalline phases (middle panel of Fig. 15.1) plus the spectrum of
Pb2+ adsorbed onto goethite (AdPb) in an inner-sphere bidentate-mononuclear
fashion, but the contribution of the AdPb spectrum was negligible in the fit.
In contrast, the dominant species of Pb in the Hamms tailings was Pb2+
sorbed to iron-(oxyhydr)oxide (48%), followed in abundance by pyromorphite
(38%) and hydrocerrusite (14%). EXAFS spectroscopy made it possible to detect
this surface-bound lead species. The presence of significant amounts of adsorbed
Pb2+ in the Hamms tailings is consistent with the near-neutral pH of associated
pore waters, and the lack of significant adsorbed Pb2+ in the Apache tailings is
consistent with the acidic pH of the pore waters, where very little Pb uptake
would be expected. The presence of surface-bound lead in the Hamms tailings
was confirmed by XPS taken before and after the SCE. Pb speciation could not
be easily determined in the combined XPS/SCE analysis because the extractions
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modified the species of Pb in both samples, showing direct evidence for redistribution of Pb in the commonly used 1 M MgCl2 extraction and successful removal
of adsorbed Pb by EDTA. The results of this study clearly show the presence of
Pb2+ adsorbed on ferric oxides in a manner very similar to that found in our model
system studies of Pb adsorption to ferric oxides (Bargar et al., 1997b).
Figure 15.1
Because of the presence of both sulphate and carbonate ligands in the two
tailings, John Ostergren also carried out combined EXAFS and ATR-FTIR studies
of the Pb/SO4/goethite and Pb/CO3/goethite systems (Ostergren et al., 2000a,b)
to determine if ternary surface complexes form in these systems. In the case of
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Ph.D. and then go to law school. He also told me that he would apply only to
Stanfords Law School and if he wasnt accepted, he would seek a position as an
environmental geochemist after completion of his Ph.D. To make a long story
short, John applied to the Stanford Law School, he completed an excellent Ph.D.
project, and he graduated from Stanford Law School (2nd in his class). John is
now an attorney with the 3M Corporation in St. Paul, Minnesota, specialising
in environmental law.
15.1.2 Lead speciation at Noyelles-Godault, Nord-Pas-de-Calais, France
Figure 15.2
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GEOCHEMICAL PERSPECTIVES | V O L U M E 1 , N U M B E R 4 A N D 5
agricultural fields used for intense farming and breeding activities. Lead levels
currently exceed 200 ppm in soils over a 40 km2 area surrounding the central
smelter. (b) Photographs of the two soil sections studied and the concentrations
of Pb and Zn as a function of soil depth (after Morin et al., 1999).
The SCE results and the Pb L III-edge EXAFS analyses of the two soils
together with spectra of several model compounds used to fit the soil spectra
(Fig. 15.3) revealed that the tilled soil contains 45% Pb sorbed on HFO, 35% Pb
associated with NOM, and 20% Pb sorbed on Mn-oxides. In contrast, Pb speciation in the wooded soil is dominated by Pb associated with NOM and Pb sorbed
on HFO. Our results for these soils with low P/Pb and Mn/Pb ratios, which were
contaminated by atmospheric deposition of PbS and PbSO4 aerosols, showed
that Pb2+ adsorbs primarily to organic matter and HFO. This joint study between
the Paris and Stanford groups (Morin et al., 1999) provided direct evidence for
adsorption processes as the main sink for Pb in organic-rich, smelter-impacted
soils.
Figure 15.3
Results of SCE of Pb from the two Pb-Zn-contaminated soils from Nord-Pasde-Calais, France and the corresponding Pb LIII -edge EXAFS spectra together
with spectra of Pb-containing model compounds used to fit the soil spectra
(T.O.C = total organic carbon; after Morin et al., 1999).
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Figure 15.4
GEOCHEMICAL PERSPECTIVES | V O L U M E 1 , N U M B E R 4 A N D 5
Figure 15.6
(a) Plot of T.O.C. and pH of the Pb-containing Largentire soil. (b) Bar graph
showing the results of the EXAFS fits of the three soil horizons (A, B, and C).
The Pb concentration data were normalised against the conservative element
Zr. Orange represents Pb2+ associated with NOM, green is Pb2+ associated with
Mn-oxides, and gray is plumbogummite (after Morin et al., 2001).
Figure 15.5
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Figure 15.7
Environmental scientists in the wild West prior to filling each other with lead.
GEOCHEMICAL PERSPECTIVES | V O L U M E 1 , N U M B E R 4 A N D 5
of arsenic associated with HFO surfaces, and in the next three subsections, we
build on our XAFS spectroscopy studies of As sorption to HFO surfaces in simple
model systems discussed in Section 11.
15.2.1 Arsenic pollution in South and Southeast Asia
Because of its toxicity and the number of people exposed to this element, arsenic
is thought to be the environmental contaminant responsible for the highest
risks of morbidity and mortality worldwide (Hopenhayn, 2006). The best-known
example of this problem is the arsenic pollution of Holocene groundwater aquifers in south and southeast Asia (India, Bangladesh, China, Myanmar, Pakistan,
Vietnam, Nepal, Cambodia), which is impacting over 100,000,000 people who
derive their drinking water from shallow, hand-drilled wells in the deltas of a
number of major rivers in this region (Brahmaputra, Ganges, Irrawaddy, Meghna,
Mekong, Red) draining from the Himalaya to the north and northwest (Charlet
and Polya, 2006; Fendorf et al., 2010) (Fig. 15.8). The redox conditions at the
depths in these deltaic sediments from which drinking water is pumped are typically anaerobic. The problem begins in the Himalaya, where abundant arsenian
pyrite is found naturally in the high-grade metamorphic and metasedimentary
rocks and ophiolites. Molecular-level studies of arsenian pyrites have shown
that significant amounts of arsenic (>1300 ppm) substitute for sulphur in pyrite
and is locally clustered in the pyrite structure (Savage et al., 2000). Oxidation of
arsenian pyrites and other As3+-bearing minerals (see ODay (2006) for a review
of common As-bearing minerals) results in the release of As3+, which is oxidised
to As5+ in the form of protonated arsenate oxoanions. These arsenate oxoanions
sorb strongly to Fe(III)-(oxyhydr)oxides (e.g., Waychunas et al., 1993; Foster et al.,
1998; Dixit and Hering, 2003; Cancs et al., 2005; Waychunas et al., 2005; Morin
and Calas, 2006; Catalano et al., 2008; Morin et al., 2008; He et al. 2010), which
can be transported in colloidal form via surface and groundwater aquifers. Such
colloids are transported by the waters of the major rivers in S and SE Asia to their
floodplains and deposited (Fig. 15.8).
There is now broad acceptance that Fe(III) and As(V) in As(V)-sorbed
Fe(III)-(oxyhydr)oxides can be bacterially reduced (e.g., Islam et al., 2004; Polizzotto et al., 2005, 2006, 2008; Kocar et al., 2006; 2008; 2009). Such reduction, as well
as reduction induced by organic carbon from human waste, in the As-polluted
deltaic sediments in the affected regions of S and SE Asia has resulted in the
release of As(III) oxoanions into groundwater, which is impacting the health
of millions of people in these regions (Fendorf et al., 2010). Thus, As-bearing
minerals, natural Earth-surface processes, and humans have conspired to create
the largest mass poisoning in human history in southern Asia. Additional details
about the environmental mineralogy of As and its impact on humans can be
found in a special issue of Elements (2006) and in Charlet et al. (2011).
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Figure 15.8
(top) Schematic drawing of the Himalaya, the Mekong River, and the Cambodian floodplain, showing the breakdown of arsenian pyrite in the Himalaya,
transport of Fe(III)-(oxyhydr)oxide colloids with sorbed As5+ (referred to as
As-Fe-hydroxide), and their deposition in the Cambodian floodplain. (lower
left) STM image of Shewanella oneidensis MR-1 on a haematite (0001) surface
(from K. Rosso, PNNL, pers. comm.). (lower right) Photo of a water well in
Cambodia that may be polluted by As3+. A highly simplified version of the
reduction reaction of Hx AsO 4 -sorbed Fe(OH)3 is shown at the bottom (from
Brown and Calas, 2011, with permission from Elsevier).
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Figure 15.9
The Vieille Usine site at Auzon (Haute-Loire, France) before (a) and after
(b) the destruction of the processing plant in 2000.
Figure 15.10
Spectacular decrease of element concentration with soil depth at the contaminated site of Auzon. (cf the log-scale for representing the concentration values;
after Cancs et al., 2005).
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Cancs et al. (2008) showed through SCE analyses that in the topsoil horizons, As is hosted by arseniosiderite (Ca 2Fe3+3(AsVO4)3O23H 2O), a secondary
mineral that forms upon oxidation of primary As-bearing minerals like arsenopyrite. At depth, more than 65% of the As is released by an oxalate extraction
step, suggesting a major association of arsenic with HFO within the soil profile,
except for the topsoil. Arsenic K-edge XANES spectra of soil samples were very
similar, suggesting that As mainly occurs mainly as sorption complexes and/or
coprecipitates. The topsoil samples have a minor amount (7%) of As3+, perhaps
related to the presence of NOM, and the spectra of the soil samples compare well
with those of As5+ sorbed or coprecipitated onto/with HFO. Linear combination
fitting of model compound spectra and shell-by-shell fitting showed that As
occurs mainly as As-bearing HFO (65%) and arseniosiderite (35%) in the topsoil
horizon (0-7 cm depth).
Similar analyses also revealed that very little arseniosiderite is present
below 15 cm depth, and that As(V) is mainly (at least 80 wt.%) associated with
HFO. EXAFS-derived As-Fe distances of 3.3 correspond to As(V) linked to
the surfaces of HFO as bidentate-binuclear complexes (Fig. 15.11), as reported
for experimental As(V) sorption on, or coprecipitation with, Fe(III)-(oxyhydr)
oxides (Waychunas et al., 1993; Foster et al., 1998; Sherman and Randall, 2003).
Schultenite, which was identified by XRD in a separate thin white layer, accounts
for < 10 wt.% of arsenic in soil samples. This study shows the importance of
bidentate-binuclear As surface complexes in the soils at this site. Arseniosiderite,
which most likely formed by oxidation of arsenopyrite, is progressively dissolved
and replaced by less soluble, poorly ordered
As-bearing HFO, which are the main hosts
for As in well-aerated soils.
Considering the decreasing solubility
of As-bearing co-precipitates as their As/Fe
ratio decreases, the high As/Fe ratio measured in the As-bearing HFO at this site
suggests that these solids could progressively transform into less soluble coprecipitates with lower As/Fe ratios. Although this
anticipated scenario is very encouraging
because it should lead to the occurrence of
less soluble As-bearing HFO, this possible Figure 15.11 Adsorption of As(V) on
transformation through a dissolutionHFO as a bidentate precipitation process poses the question of
binuclear complex.
the possible long-term mobility of As during
such a transformation. However, the efficiency of such processes that limit arsenic
migration cannot be estimated because past fluxes of arsenic are unknown.
Arsenic speciation in soils developed over a geochemical anomaly: Changes in
arsenic mineralogy over geological time periods. Arsenic geochemical anomalies are
widely used in geochemical exploration for Au, W, Sn, and other ore deposits.
Arsenic geochemical anomalies also provide examples of long-term regional
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Figure 15.12
642
As speciation in the soil profile at Echassires. (a) Photo of the soil profile. (b)
As, Cu, and Zn concentrations as a function of soil depth. (c) XRD patterns of
soils from different depths (primary pharmacosiderite = Ph; muscovite = Mu;
biotite = Bt; after Morin et al., 2002).
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Figure 15.13
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surfaces in the topsoil. This interpretation of the EXAFS spectra of the soils
explains the association of As with HFO and is consistent with SEM-EDS and
electron microprobe analyses of the soils. Soil evolution leads to the remobilisation of As(V), probably after partial dissolution of pharmacosiderite, with
As(V)-oxoanions being eventually held as inner-sphere complexes at the surface
of Fe-(oxyhydr)oxides. In this case, adsorption provides long-term trapping of the
arsenic, given the slightly acidic and oxidising conditions, and delays potential
release of As into the biosphere.
15.2.3 A natural bioremediation site at the Carnouls,
Gard, France acid mine drainage
A great example of natural attenuation via the formation of As-Fe mineral
phases is provided by an AMD system with exceptionally high concentrations
of dissolved arsenite, located at Carnouls, Gard, France (Fig. 15.14). This study
started after I (GC) gave a seminar in 1999 at the University of Montpellier on
environmental mineralogy, where I discussed with Marc Leblanc, a French ore
geologist, whom I knew since the early 80s, about his recent observations on the
evolution of mine tailings at Carnouls (Leblanc et al., 1996). Guillaume Morin
coordinated at LMCP/IMPMC an in-depth study of what appears to be, 10 years
later, one of the best-known examples of an As-contaminated AMD. This AMD
system is generated by 1.5 million tons of mine tailings containing arsenious
pyrite, associated with Pb-Zn mineralisation in Triassic sandstones that was
mined for a few decades. Below the Carnouls tailings impoundment, Reigous
Creek (Fig. 15.14) supplies high As concentrations, as soluble (up to ~4 mgl1)
and particulate (up to 150 mgAs g1) forms to the Amous River, located within
the drainage basin of the Rhne River that eventually reaches the Mediterranean Sea. Biomineralisation (bacterial stromatolite and bio-sediments) at the
Carnouls AMD site results in a remarkable accumulation and concentration of
As in a mineral form via direct or indirect microbial action. This biomineralisation
limits and controls As-pollution downstream. Here, bacteria play an important
role, and this biogeochemical coupling causes spatial and seasonal modifications
in As speciation (Morin et al., 2003; Casiot et al., 2005). Ferrihydrite and poorly
crystalline Fe-(oxyhydr)oxides are present in the Carnouls AMD system and
are common, in general, in AMD systems. Under low pH and oxidising conditions, mineral oxidation is enhanced by the metabolic activity of bacteria such as
Acidithiobacillus ferrooxidans, which catalyses the oxidation of Fe(II) to Fe(III) by
dissolved O2 and leads to the formation of ferrihydrite and other poorly crystalline Fe-(oxyhydr)oxides or oxysulphate phases.
At Carnouls, a strain of A. ferrooxidans promotes the formation of tooeleite,
a ferric arsenite mineral (Fe6(AsO3)4SO4(OH)44H 2O) identified as the main
constituent of the stromatolite-like deposits shown in Figure 15.14a (Morin et
al., 2003, 2007). The formation of tooeleite results in a dramatic decrease in the
concentration of As(III) in Reigous Creek.
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Figure 15.14
(a) The AMD site of Carnouls. (b) Reigous Creek merges with the Amous River.
The pH increases, reducing the As(III) concentration. Note that the mine site is
inhabited. (c) Map showing Reigous Creek which has a low pH and before and
at the stromatolite has exceptionally high As(III) concentrations, which decrease
due to bacterial activity near the stromatolite (after Morin et al., 2003).
At or near the source of AMD, the anoxic and acid (pH = 2-3) waters are
strongly enriched in Fe (0.5-1 gl1), As (50-350 mgl1), and sulphate (1-3 gl1).
Under such conditions, arsenite is the predominant dissolved arsenic species.
After 30 m of downflow, 20-60% of the As has been removed via the biogenic
precipitation of rare mineral species, such as amorphous ferric arsenite hydroxysulphates and nanocrystalline tooeleite. This intense microbial activity results
in an efficient detoxification of the AMD waters although the As-rich mineral
phases formed in the upstream zone are fairly soluble. Further downstream, the
progressive decrease of the dissolved arsenic concentration allows the precipitation of schwertmannite, which traps most of the remaining dissolved arsenic.
The dissolved arsenic concentration thus decreases to reach about 1 mgl1 before
the confluence with the non-polluted river, 1.5 km downstream from the waste
pile (Casiot et al., 2005).
An important finding of these studies concerns the influence of bacterial
activity on the nature of the minerals forming at this site. Arsenic K-edge XANES
spectra indicate a seasonal variation of As speciation in the acidic spring zone
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hydroxysulphates with As:Fe mole ratios up to 0.7 (Leblanc et al., 1996; Morin
et al., 2003). This value approaches the upper limits observed for synthetic coprecipitates of As(V)-HFO phases (Carlson et al., 2002). EXAFS analysis of these
mineral phases confirms the similarity between the natural and in vitro mineral
phases. In addition, spectroscopic evidence of the formation of stable arsenic
complexes on the schwertmannite surface (Waychunas et al., 1995) explains the
inhibiting role of As in the crystallisation of schwertmannite, leading to the
formation of amorphous As-rich ferric hydroxysulphates (Carlson et al., 2002).
Figure 15.16
Figure 15.17
(a) As(III)-oxidising Thiomonas sp. (yellow arrows), which are active during
spring/summer, promote the formation of amorphous Fe(III)-As(V) hydrous
oxides. (b) As(V) inner-sphere complexes on schwertmannite (after Morin
etal., 2003).
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Figure 15.18
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WEATHERING AS A MOLECULAR-LEVEL
16. MINERAL
MINERAL-SURFACE PROCESS
When rocks are exposed at Earths surface, their equilibrium is disturbed and
their minerals react and experience transformations, resulting in the formation
of soils and the release of ions, which modify the composition of ground and
surface waters. The rates of mineral weathering reactions depend on many factors,
including solution composition, crystal structures, bond energies, and ionic transport from reaction sites, resulting in kinetic rather than purely thermodynamic
controls. Proton-promoted and ligand-catalysed mineral dissolution are key
weathering processes occurring at mineral-water interfaces, as shown by Stumms
innovative approach to mineral weathering based on mineral-surface coordination chemistry (Furrer and Stumm, 1986; Zinder et al., 1986; Biber et al., 1994).
These enlightening studies demonstrated the links between molecular structure
and interfacial reactivity and resulted in the widely used surface-coordination
model, which describes mineral surfaces as comprised of distinct reactive sites.
Coordination changes (e.g., for Al, Fe) and redox changes (e.g., for Mn, Fe)
are major transformations observed during mineral weathering and soil formation. Phases resulting from low-temperature weathering provide important information on the history of continental surfaces and their evolution as a result of
human activities or climatic forcing conditions. The high specific surface areas
of these phases explain their high surface reactivity, which drives key processes
such as trace element adsorption and incorporation, crystal growth, and phase
dissolution.
Although neither of us has been directly involved in experimental studies
of the dissolution of major rock-forming minerals, we have included this topic
because of the importance of dissolution reactions of feldspars and other major
rock-forming minerals in (1) the cycling of elements in the biosphere, (2) controlling atmospheric CO2 levels, (3) controlling the composition of natural waters,
(4) the formation of soils, (5) maintaining climatic stability over eons, and
(6) causing climatic swings in response to tectonic and paleogeographic factors
(e.g., Kump et al., 2000). In spite of their importance, however, mineral dissolution
reactions are still not well understood at a fundamental, mechanistic level under
equilibrium or far-from-equilibrium conditions. Moreover, the major difference
in estimated rates of mineral-weathering reactions in the field vs. those measured
in the laboratory underlines our current state of understanding of the factors
controlling mineral dissolution in complex, natural environments (e.g., White and
Brantley, 1995), although new understanding has emerged (Nugent et al., 1998;
White and Brantley, 2003; Maher et al., 2006; Maher, 2010).
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Figure 16.1
(a) In situ X-ray reflectivity vs. time (measured at the anti-Bragg condition,
shown in the inset at top) during dissolution of orthoclase feldspar, KAlSi3O8 ,
(001) cleavage surface at extreme pH values. The removal of successive monolayers (ML) is noted for each set of data. (b) in situ CTR diffraction profiles for
a freshly cleaved orthoclase (001) surface (circles) and after reaction at pH = 2.0
(1 and 15 ML dissolved; diamond and square) and pH = 12.9 (2 ML dissolved;
triangle; after Teng et al., 2001).
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Figure 16.2
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655
[ Zr ]rock
MTCx = [ X ]rock [ X ]soil
[ Zr ]soil
/ (100 m x )
(16.1)
where [X]rock and [X]soil are the concentration of element x (not the oxide) in the
bedrock and in the soil (both values in wt.%), respectively, [Zr]rock and [Zr]soil are
the concentration of Zr in rock and soil, respectively, and m x is the molar mass
of element x. As a component of resistant primary minerals and a low solubility
element in aqueous solutions, a geochemical invariant must be homogeneously
distributed in the parent rock. Ti and Zr are often used as invariants, but this
has been questioned relative to more immobile elements such as Th (e.g., Braun
et al., 1993).
16.2.2 Surface chemistry of zircon in soils and sediments
The resistance of zircon to dissolution is exceptional, as compared to other silicates. Figure 16.3 shows the preservation of a detrital zircon sampled in laterites developed on deeply weathered continental sediments from the Amazon
Basin, hence having suffered several alteration and weathering stages, including
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metamictisation a process that results in structural damage and loss of longrange order in zircons because of the recoil of included radioactive element nuclei
(e.g., U and Th) as the result of their alpha particle emissions. Despite this complex
geological history of zircons in sediments and soils, a pristine shape is well
preserved.
Figure 16.3
The presence of a Zr-rich protective layer at the zircon surface has been
ruled out using Rutherford backscattering (RBS) and XPS, which sample depths
of ~10 nm and <m, respectively (Balan et al., 2001c). Indeed a mineral surface
corresponds to a volume, as it extends from the top atomic layer to a few m deep,
depending of the technique used. 4He+ RBS data on well-crystallised zircons from
soils and underlying Cenozoic sediments in the Manaus region (Brazil) (Fig. 16.4)
show signals related to O, Si, Zr, and H in the upper 20 nm and the absence of
Zr enrichment at the zircon surface. Zr 3d XPS spectra show only a contribution
from zircon, without any surface species such as ZrO2 of Zr(OH)4. These data
confirm the absence of significant chemical change at the zircon surface during
or following weathering.
The absence of Zr enrichment at the surface of zircon grains is observed
for all depths investigated. Even when dissolution features are observed (see
Fig. 16.3b), the slow dissolution rate of zircon prevents any dissolved Zr from
reaching saturation concentration, despite the low solubility of Zr oxides. This
low dissolution kinetics is the principal factor limiting Zr mobility in the investigated Amazonian soils. These findings demonstrate that zircon dissolution
during weathering is controlled by surface reactions, as is the case for other
orthosilicates (Brady and House, 1996). The absence of a protective passivation
layer of Zr oxides/hydroxides is consistent with the PillingBedworth rule used
in corrosion science: if the volume of the product is greater than the volume of the
reactant, a protective surface layer may have formed during weathering (Velbel, 1993).
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Here, the ratio of the molar volume of the zirconium oxide (i.e. baddeleyite) to that
of zircon is 0.52, i.e. significantly smaller than 1, and a passivating baddeleyite
layer is not predicted.
Figure 16.4
(a) 4He+ RBS spectra of a well-crystallised zircon from a soil near Manaus with
data shown as red symbols, the simulation as a bold line, and the element
contributions as the thin lines. (b) Zr spectral region of the same sample
compared to a partially metamict zircon from the same soil (red and black
circles, respectively). The metamict zircon shows some Zr depletion as a result
of the presence of kaolinite contamination near the surface and the calculated
influence of a 0 to 20 nm ZrO2 surface layer is shown by the solid lines (after
Balan et al., 2001c).
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Figure 16.8
Measured vs. calculated radiation dose suffered by zircons from the Manaus
region. The thin dashed line corresponds to a damage of ~3.5 1015 -events
mg 1. The shaded zone corresponds to the estimated position of the first
percolation threshold at 3 4 1015-events mg 1 (after Balan et al., 2001b).
Perspectives on Mineral Weathering We have touched only briefly on mineralwater interfaces and their effect on weathering processes and have not addressed
the effect of biology despite the large emerging literature on this topic. That would
be an excellent topic for a future Geochemical Perspective issue. Nevertheless, Werner
Stumm and his co-workers developed a very useful non-biotic view of mineral weathering based on a surface coordination-chemistry approach that has proven very
useful and those early results are still consistent with more recent studies discussed
above. The Stumm approach was based on the idea that mineral dissolution involves
surface chemical reactions with H+, OH , metal ions, and ligands to form an array of
surface complexes whose reactivities determine the mechanisms of many surface-controlled
processes. The pH-dependent charging behaviour of the mineral surface plays a key
role in attracting ions of opposite charge and thus helps control weathering rate.
Stumm and co-workers extended the model to include the effects of inhibition such
as those due to the formation of binuclear surface complexes with oxoanions such
as phosphate, arsenate, and sulphate, which were found to inhibit reductive as well
as nonreductive dissolution of metal oxides. They attributed this effect to the large
activation energy that must be overcome to detach two metal surface centres and
the lack of additional surface protonation when uncharged binuclear or multinuclear
complexes are formed. In addition, they suggested that multivalent cations such
as Al3+ are effective inhibitors of dissolution because they block surface sites and
decrease surface protonation, particularly in acidic solutions. This mechanism, as
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well as other inhibitory mechanisms involving organic ligands, are consistent with
the suggestion by Maher et al. (2006) that the lower observed rate of silicate mineral
dissolution in the laboratory vs. in the field (the so-called dissolution rate conundrum) is largely the result of the gradual loss of reactive sites on silicate surfaces with
time. This suggestion is also consistent with the conclusions of White and Brantley
(2003) that progressive depletion of reactive surface sites with weathering, resulting
in part from accumulation of secondary surface precipitates, explains the difference
in weathering rates of fresh and weathered plagioclase feldspar under otherwise
identical experimental conditions. Geochemists still dont have good experimental
procedures for estimating reactive surface areas of minerals undergoing weathering
under equilibrium or far from equilibrium conditions. White and Brantley also point
out that experimental dissolution rates are measured at high fluid/mineral ratios over
short time periods, which contrasts with natural weathering where fluid-mineral
ratios are lower and reaction times are much longer conditions that cannot be easily
duplicated in the laboratory. As a result, we have some way to go before accurate
quantitative models of mineral weathering under various conditions are available.
Such models will require information about mineral-surface dissolution mechanisms,
which is mostly lacking.
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Figure 17.1
Biogeochemical cycle of aluminium including some clues about its molecularscale environment (modified from Tria etal., 2007).
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665
Figure 17.2
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Figure 18.1
(18.1)
Mn4+
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stabilisation energy difference (228 kJmol1) between these two oxidation states
drives the redox equilibria towards oxidised Co species at Mn-oxide surfaces and
explains the spectacular trapping efficiency of Co by Mn oxides.
Figure 18.2
Crystal-field splitting of 3d-orbitals: for the octahedral Co2+ (high spin), the
usual oxidation state of Co in minerals. Under highly oxidising conditions, such
as those prevailing in the Mn-oxides occurring in laterites, Co3+ prevails. Crystalfield theory predicts that a low-spin configuration is the most stable, due to
the presence of the six 3d electrons in the stabilising xy, yz and zx orbitals.
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Figure 19.1
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671
complex on the surfaces of experimental and natural samples. The local coordination environment of Mo in this polynuclear complex is a distorted octahedron,
whereas Mo in solution is predominantly in the form of MoO42. The calculated
fractionations for MoO42 versus Mo6O192 fit the experimental data of Wasylenki
et al. (2008) very well (Fig. 19.2), and their results indicate that the difference in
coordination environment between dissolved Mo and adsorbed Mo is the cause
of 97/95Mo isotope fractionation.
Figure 19.2
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supported by the Stanford Global Climate and Energy Project (GCEP). Our
approach is multifaceted and involves (1) experimental studies of mineral-fluid
reactions at temperatures and pressures relevant to CO2 storage that utilises
Dickson-type rocker bombs, (2) detailed characterisation of reaction products
using spectroscopic and X-ray scattering methods (XPS, XRD) and nm-scale
imaging (STXM, SEM, TEM), (3) isotopic tracer measurements coupled with
secondary ion mass spectrometry (SIMS) to track chemical and physical processes
occurring in the 100 to 1000 nanometre thick interface between the fluid and the
unreacted Mg-silicate crystal, (4) thermodynamic modelling of the carbonation
reactions, and (5) field studies of a natural analogue of mineral carbonation of
mafic and ultramafic rocks at Red Mountain, California, USA. Our carbonation
experiments using forsterite (Mg2SiO4) have shown that (1) reaction rates are
enhanced in the presence of supercritical CO2 relative to pure water at the same
temperature and pH, (2) olivine dissolution rates decline as the experiments
approach equilibrium with amorphous silica, likely due to formation of a Si-rich
layer on the olivine surface, and (3) the Si-rich coating on forsterite is more
consistent with a leached layer than a re-precipitate, based on measurement of a
29
Si label using SIMS, which challenges the generality of the universal dissolution and reprecipitation model of silicates proposed by Hellmann et al. (2003,
2011, 2012), Daval et al. (2011), and King et al. (2011). Knowledge of the mechanism of formation of the apparent silica passivation layer is of critical importance
in devising ways to prevent its formation and its effect on olivine dissolution rate.
This is currently one of the main problems in developing an efficient and costeffective process for carbonate mineralisation of Mg-silicates.
Perspectives on Mineral Surface-Controlled Isotope Fractionation and MineralCarbonation Reactions Although very little detail was presented in the previous
two short sections, hopefully, it is clear from the examples discussed that mineralsurface reactions play a major role in the realm of isotope geochemistry and in controlling the geochemical portion of the carbon cycle. This note was added to remind the
reader that understanding the molecular-level processes controlling isotope fractionation and the reaction of CO2 with mineral surfaces helps us understand how
environmental variables affect these processes at a fundamental level.
674
GEOCHEMICAL PERSPECTIVES | V O L U M E 1 , N U M B E R 4 A N D 5
The multiple
ba r r ier d i sposa l
concept, used to
contain radioactive
elements from highlevel and long-lived,
intermediate-level
wastes, is based on
(1) durable matrices,
(2) stainless steel
containers, (3) engineered ba r r iers
(bentonitic clays and
concrete) = the near
field, and (4) the
geological surroundings = the far field.
At all stages, molecular-scale reactions
at mineral(glass)w ate r i nte r f ac e s
are key processes
governing nuclear
w a ste b eh av iou r
over the long term.
Waste forms must
have chemical and
mechanical durabilit y against the
Figure 21.1
GEOCHEMICAL PERSPECTIVES | G O R D O N E . B R O W N J R . G E O R G E S C A L A S
675
forcing conditions represented by chemical alteration and irradiation. Vitrification of high-level radioactive waste is currently used on an industrial scale in
several countries (Grambow, 2006). Glass durability is thus a major issue in the
perspective of a deep geological disposal due to the need to model the long-term
behaviour of nuclear glasses. For instance, the formation of an alteration layer
(=gel) during alteration influences radionuclide retention (Frugier et al., 2008).
Molecular-scale approaches provide a unifying view of the processes operating
at glass-water interfaces.
Figure 21.2
676
GEOCHEMICAL PERSPECTIVES | V O L U M E 1 , N U M B E R 4 A N D 5
Figure 21.3
(a) Principal glass alteration mechanisms. (b) Glass alteration kinetics versus
time for an R7T7-type glass in a static system at 90 C (from Vernaz, 2002,
with permission from Elsevier Masson SAS).
GEOCHEMICAL PERSPECTIVES | G O R D O N E . B R O W N J R . G E O R G E S C A L A S
677
Figure 21.4
678
GEOCHEMICAL PERSPECTIVES | V O L U M E 1 , N U M B E R 4 A N D 5
Fe K-edge GI-XAFS (see Section 5) probes the evolution of the local environment around Fe at the glass surface. Both GI-XANES and GI-EXAFS spectra
indicated a transformation of IV Fe3+ in the pristine glass to VI Fe3+ in the alteration
layer (Fig. 21.5). The transformation was complete after 1 day. Fe-O distances
increased from 1.88 to 2.03 between bulk and surface. The local structural
environment of VI Fe3+ mimics that of HFO (Combes et al., 1989; Allard et al., 1999;
Cismasu et al., 2011). The Fe-Si contribution observed in the glass at 3.00.04
was replaced at the glass surface by Fe-Fe contributions at 3.12 and 3.42 ,
characteristic of HFO (Combes et al., 1989; Allard et al., 1999). Complementary
Fe-LII,III XANES spectra showed that the IV Fe3+ to VI Fe3+ coordination change
starts after a few hours of alteration, irrespective of the alteration conditions.
Figure 21.5
(21.1)
Na+
in which
sol represents the sodium released into the solution. Hydration of
alkalis explains their decreasing concentration near the glass surface. During
hydration, alkalis lose their charge-compensating role, causing an underbonding
of the oxygens bonding Si tetrahedra and cationic sites, and eventually forcing
cations to adopt a different local topology. For Fe3+, this topological change corresponds to the precipitation of HFO at the glass surface:
Na(FeO2)+ H+ (FeOOH) + Na+
GEOCHEMICAL PERSPECTIVES | G O R D O N E . B R O W N J R . G E O R G E S C A L A S
(21.2)
679
in which (FeO2) represents IV Fe3+ in the starting glass. The presence of HFO is
observed during the alteration of basaltic glasses (Toner et al., 2012) and is also
predicted by geochemical codes such as DISSOL (Crovisier et al., 1992). The structural evolution around Fe3+ during glass alteration corresponds to a dissolutionprecipitation process at near- and under-saturated conditions.
Zirconium, an element of great technological importance, is an interesting
structural probe, due to its versatile geometry, including coordination numbers
ranging from 6 to 8. Electron detection mode Zr-LII,III XANES data (Fig. 21.6)
indicate that at near-saturated conditions Zr remains as VI Zr, as in the pristine
glass. In contrast, at under-saturated conditions, Zr coordination changes to VII Zr
as in hydrous zirconia (HZO), Zr(OH)4. As alkalis/alkaline earths are implied in
local charge compensation of VI Zr , this coordination change reflects modifications in the structure and composition of hydrated glasses. The structural evolution of Zr depends on alteration conditions because of its ability to be charge
compensated by cations such as Ca 2+, which is still observed in the gel layer. In
contrast, IV Fe3+ ions, which are mostly charge compensated by alkalis, are leached
in the alteration solution.
Figure 21.6
680
Zr-LII,III XANES spectra of reference samples and altered glasses. (a) Crystalline
references: VIZr (lemoynite), VIIZr (baddeleyite, and HZO) and VIIIZr (zircon),
showing a major influence of Zr-coordination number and site geometry on
the spectra. (b) Evolution of the surface of glass monoliths altered in undersaturated and near-saturated conditions at 90 C (after Plegrin et al., 2010).
GEOCHEMICAL PERSPECTIVES | V O L U M E 1 , N U M B E R 4 A N D 5
Glass alteration at under-saturated conditions corresponds to a dissolutionprecipitation process. The Zr-coordination change is completed after a few hours,
i.e. much earlier than the appearance of the saturation regime. This may be
explained by the small thickness probed by surface-sensitive XAFS. Leaching
progresses inward, with the external part of the alteration layer being first aged
during alteration. The lack of charge-compensating cations and protonation of
the oxygen bonded to Zr and Si atoms results in the breakage of bonds between
Zr sites and the silicate glassy network, due to the over-bonding of the oxygen
atoms, giving rise to an HZO local structure. Under these conditions, the gel is
depleted in Ca, which is not needed for charge-compensation.
Near-saturated conditions correspond to an in situ condensation process.
The residual alteration comes more from diffusion through the altered surface
than from a modification of the near-surface. Under these conditions, Zr retains
the same coordination number and linkages with the silicate framework as in
the glass. Ca plays the role of an alternate charge-compensating cation during
alteration, as alkalis are leached. This is at the origin of lower leaching rates and
higher glass durability on the short term.
A sketch of the molecular-scale processes during glass alteration in undersaturated conditions is presented in Figure 21.7. These diagrams illustrate the
general use of Paulings rules for predicting the links between the polyhedra in
complex materials such as glasses and gels.
Figure 21.7
GEOCHEMICAL PERSPECTIVES | G O R D O N E . B R O W N J R . G E O R G E S C A L A S
681
Perspectives on the Long-Term Behaviour of Glass Surfaces New glasses are needed
for the waste generated by nuclear mixed oxide fuel and other alternate fuels, as well
as by future generations of nuclear reactors, such as 4th generation reactors (GEN IV).
A major question concerns the chemical dependence of both the initial alteration rate
and the residual rate over the long term. This positive structure-property relationship has been discussed above as arising from the connectivity between the ZrO6
octahedra and the silicate network of the alteration gel. Experimental studies and
Monte Carlo simulations of the mesoscopic structure of the alteration gel show that
glasses with high dissolution rates undergo fast restructuring and corrode slightly,
and that glasses with low dissolution rates undergo slow restructuring and corrode
deeply (Cailleteau et al., 2011). All of these molecular, and mesoscopic-scale observations, result in specific implications for the long-term behaviour of nuclear glasses
under geological repository conditions. The long-term stability of a glass waste form
depends not only on the leaching processes at the glass-solution interface, but also
on the long-term stability of the alteration layers.
682
GEOCHEMICAL PERSPECTIVES | V O L U M E 1 , N U M B E R 4 A N D 5
GEOCHEMICAL PERSPECTIVES | G O R D O N E . B R O W N J R . G E O R G E S C A L A S
683
Figure 22.1
We have learned a great deal about molecular-level processes occurring at solid-water interfaces since the pioneering studies of solidwater and solid-gas interfaces by Irving Langmuir and of mineral-water
interface chemistry by Werner Stumm and Paul Schindler and their
students. Konrad Krauskopf also deserves special mention because of his
pioneering study of the factors that control the trace element composition of seawater, with adsorption of trace elements on mineral surfaces
exerting a primary control (Section 4).
The electrical double layer (EDL) model of solid-water interfaces developed by Helmholtz, Gouy, Chapman, Stern, and Grahame in the late 19th
and early 20th centuries to explain the surface-charging behaviour and
stability of colloidal suspensions in water has been shown to be valid
qualitatively by modern molecular-level, synchrotron-based studies and
by the applications of ab initio MD simulations and DFT (Section 7.4).
Our understanding of the structure of the EDL has advanced tremendously over the past thirty years through the application of a variety
of in situ synchrotron-based methods. These methods have provided
684
GEOCHEMICAL PERSPECTIVES | V O L U M E 1 , N U M B E R 4 A N D 5
GEOCHEMICAL PERSPECTIVES | G O R D O N E . B R O W N J R . G E O R G E S C A L A S
685
686
GEOCHEMICAL PERSPECTIVES | V O L U M E 1 , N U M B E R 4 A N D 5
GEOCHEMICAL PERSPECTIVES | G O R D O N E . B R O W N J R . G E O R G E S C A L A S
687
Multiple analytical and experimental methods, including non-synchrotron methods, should be used in characterising mineral-water interfaces. (e.g., EXAFS spectroscopy should more often be complemented
by SCE to help determine the speciation of surface species)
Consideration of potential sample damage by intense synchrotron
radiation beams or electron imaging is essential. For example, intense
synchrotron X-ray beams at the APS converted Cr(VI) in Hanford, WA
drill core samples to Cr(III) (e.g., Zachara et al., 2004).
In situ methods (i.e. methods allowing water and/or gases to be present
and in contact with sorbent surfaces under ambient P-T conditions)
should be used where possible to achieve conditions as close as possible
to those occurring in the natural environment. This point is related to
the pressure gap mentioned above.
Theoretical modelling must be done on realistic systems (e.g., dont
expect the reaction of one water molecule per surface unit cell to realistically model the reaction of water with a mineral surface). Ab initio
thermodynamics is a good example of a theoretical approach that simulates realistic P and T values and gases in contact with mineral surfaces.
MD simulations are only as good as the interatomic potentials used.
A potential for water that does not allow water to dissociate does not
capture one of the important properties of liquid water that occurs at
mineral-water interfaces.
Theoretical methods can provide unique insights about H-bonding at
mineral-water interfaces; currently, there are no experimental methods
that can provide this information.
688
GEOCHEMICAL PERSPECTIVES | V O L U M E 1 , N U M B E R 4 A N D 5
Studies should be conducted on the appropriate environmental interfaces for a given problem. For example, reductants or oxidants that
control the kinetics of geochemical interface reactions, such as the
reduction of Cr(VI) to Cr(III), should be chosen rather than reductants
or oxidants that result in very slow kinetics. In the case of Cr(VI) reduction, aqueous Fe(II) or S(-II) are far more effective reductants of Cr(VI)
to Cr(III) than NOM (Fendorf et al., 2000).
GEOCHEMICAL PERSPECTIVES | G O R D O N E . B R O W N J R . G E O R G E S C A L A S
689
690
GEOCHEMICAL PERSPECTIVES | V O L U M E 1 , N U M B E R 4 A N D 5
GEOCHEMICAL PERSPECTIVES | G O R D O N E . B R O W N J R . G E O R G E S C A L A S
691
and NOx to Earths troposphere during the period since the industrial revolution.
The effect was the production of significant acid rain in areas such as northern
Europe and the northeastern US, which has resulted in major damage to forested
ecosystems. The second is the use of insecticides such as DDT (dichlorodiphenyltrichloroethane) to control malaria-carrying mosquitoes. As a child growing up
in Mississippi, I (GB) also played in the DDT clouds emitted by trucks traveling
through residential neighbourhoods in the summer months. At the same time and
for the same reason,
DDT wa s spread
by helicopters on
brack ish la kes in
Southern France near
the summerhouse of
my (GC) family. The
long-term effect of
DDT was the death
of birds, as described
in Rachel Carsons
famous 1962 book
entitled Silent Spring,
which helped initiate
the env ironmental
movement in the
US and Europe. The
third example is the
use of laundry detergents, which contain
ph o s ph a t e s , a n d
fertilisers, which in
turn contain nitrates
and other chemicals.
The release of such
chemicals into waste Figure 22.2 Cause-and-effect relationships between anthropo- water that eventugenic environmental perturbations and the effects
ally reaches surface
of such perturbations. The lag time between the
perturbation and the effect is often 20 to 30 years. waters such as lakes
in Switzerland, the
Mississippi River, and portions of the Gulf of Mexico has resulted in major
eutrophication of the waters. For example, in the Gulf of Mexico just south of
New Orleans, these chemicals have resulted in a major dead zone in which
shellfish such as shrimp and oysters have been impacted. In the last case, which
shows a fullerene molecule as an example of an engineered nanoparticle, we dont
yet know the environmental consequences of the growing variety of engineered
and incidental nanoparticles being introduced into the environment. The question is how long will it take us to recognise the effects of these nanoparticles on
the health of organisms, including humans, and ecosystems.
692
GEOCHEMICAL PERSPECTIVES | V O L U M E 1 , N U M B E R 4 A N D 5
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693
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732
Geochemical Perspectives | V o l u m e 1 , N u m b e r 4 a n d 5
LIST OF ACRONYMS
AFM
Ag-NP
Silver Nanoparticle
AGU
ALS
AMD
ANDRA
APS
Brunauer-Emmett-Teller
BL
Beamline
BO
Bridging Oxygen
BSi
Biogenic Silica
Tridentate Corner-Sharing
CARS
CNRS
CRAEMS
CTR
DFT
GEOCHEMICAL PERSPECTIVES | G O R D O N E . B R O W N J R . G E O R G E S C A L A S
733
EAWAG
EDL
EDXS
EGB
EM
Electrophoretic Mobility
EMSI
ENS
EPS
EPSCI
ESRF
FEFF
F Effective
Fh
Ferrihydrite
FT
Fourier Transform
FY
Fluorescence Yield
GB
Gordon Brown
GC
Georges Calas
GCEP
GE
GI
Grazing Incidence
GI-XAFS
HFO
HRTEM
HZO
Hydrous Zirconia
IHP
734
GEOCHEMICAL PERSPECTIVES | V O L U M E 1 , N U M B E R 4 A N D 5
IMPMC
IR
Infrared
IRD
IUF
Langmuir (= 10-6Torr-sec)
LBNL
LCLS
LMCP
LMU
LPS
Low Spin
LURE
MCL
MD
Molecular Dynamics
MES
ML
Monolayer
MO
Molecular Orbital
MS
Multiple Scattering
MSA
MUSIC
NEXAFS
Near Edge X-ray Absorption Fine Structure (the term used for soft X-ray/
VUV near edge X-ray absorption spectra)
NICA
NIRT
NIST
NMR
NOM
NSF
OHP
PAA
Polyacrylic Acid
pH PZC
PNNL
PPHA
GEOCHEMICAL PERSPECTIVES | G O R D O N E . B R O W N J R . G E O R G E S C A L A S
735
PVP
Polyvynil Pyrrolidone
RAXR
RBS
Rutherford Backscattering
RH
Relative Humidity
SEM
SCE
SCM
SFG
SIMS
SLAC
SRFA
SSRL
STXM
TEM
TLM
TO
Terminal Oxygen
T.O.C.
TPD
XAFS
XANES
X-ray Absorption Near Edge Structure (the term typically used for hard X-ray
(> 4 keV) near edge x-ray absorption spectra)
XAS
XPS
XRD
X-ray Diffraction
XSW
XSW-FY
UHV
UPD
UPMC
m-SXRF
m-XANES
-XRF
v.u.
Valence Units
VUV
Vacuum Ultraviolet
Atomic Number
736
GEOCHEMICAL PERSPECTIVES | V O L U M E 1 , N U M B E R 4 A N D 5
INDEX
A
ab initio thermodynamics 540, 576, 688
acid-base chemistry 484, 545
acid mine drainage (AMD) 616, 636
Acidothiobacillus ferrooxidans 646
adsorption isotherm 520, 533, 542, 595,
685
adsorption of trace elements 663, 684
Advanced Light Source (ALS) 495
Advanced Photon Source (APS) 495, 733
AFM 493, 558, 651-653, 686, 733
Al 2O3 497, 535, 541, 542, 546, 551, 558,
566, 571, 572, 575, 576, 579, 581, 583,
587-590, 592, 596, 598-600, 602, 604,
606-611, 613, 695, 697, 703, 706-708,
710, 712, 714, 715, 717, 719, 727-729
Al K-edge XAFS 665
Allard, Thierry 505, 507, 508, 623, 624,
626, 679, 695, 696, 697, 702, 708, 723
allophane 505, 626, 712
alteration of nuclear glasses 675, 677
Amazon River 623, 656, 695-697, 708, 723
arsenate 535, 592, 619, 637, 649, 661, 699,
703, 715, 720, 729, 730
B
bacteria 603, 606, 707, 717
Balan, Etienne 506, 626, 656-658, 660,
661, 695, 696, 699, 705, 707, 708
Beevers, C. Arnold 513, 514
bioavailability 493, 602, 609, 610, 628, 629,
632, 649, 667, 685, 686, 701, 706, 721
biofilms 496, 545, 554, 603, 604, 607, 609,
610, 616, 685, 687, 704, 707, 718, 726,
727
biogenic silica (BSi) 663, 666, 705, 709,
715, 728, 729, 733
Bird, Dennis 499, 522, 673, 715, 724
boehmite 566, 596-601, 653, 731, 732
GEOCHEMICAL PERSPECTIVES | G O R D O N E . B R O W N J R . G E O R G E S C A L A S
737
C
CaO 566, 567, 569, 572, 583, 717
Carnouls 644-646, 649, 720
cation exchange 546
CD-MUSIC model 553, 598
Chaka, Anne 497, 498, 577, 709, 717-719,
727-729
chemisorption 571, 697, 700
chromium 707, 732
Cismasu, Cristina 499, 613, 616, 617, 619,
626, 679, 703, 704, 719
cleanliness gap 687
CO2 485, 491, 499, 631, 650, 663, 673-675,
683, 690, 697, 700, 705, 714, 715, 717,
719, 721, 722, 725, 732
cobalt 484, 667, 668, 701, 705, 714, 716,
718, 721, 727
coesite 517, 518, 719
colloidal particles 548, 560, 622, 623
Combes, Jean-Marie 492, 503, 615, 626,
679, 704, 718
Constant Capacitance model 552
continuous random network 561, 689
coordination number 491, 509, 510, 513,
516, 530, 532, 534, 586, 659, 660, 668,
680, 681
crystal-field 502, 667
Crystal-field splitting 668, 669
crystal growth 650
crystal truncation rod (CTR) 497
D
Debye-Waller factors 532
Density functional theory (DFT) 540, 564,
695, 696, 712, 715, 717, 718, 730, 732,
733
diatom 506, 663
diatoms 663, 664, 665
dielectric constant 543, 550, 551, 553, 556,
558, 560, 581, 583, 584, 689, 727
diffuse double layer model 552, 719
738
E
effective ionic radii 509, 516, 725
electrical double layer (EDL) 484, 496, 684
electrophoretic mobility 543
electrophoretic Mobility 734
electrostatic interaction 550, 560
Ewing, Rod 507, 660, 675, 706, 724, 730
EXAFS 529-533, 536-538, 555, 589-594,
605, 615, 619, 626, 629-631, 633, 634,
641, 643, 644, 647, 660, 664, 667, 668,
676, 678, 679, 688, 697, 700-702, 704,
705, 710, 717-721, 723, 725, 726, 729,
730, 734
F
Farges, Francois 492, 497, 503, 507, 532,
660, 676, 677, 698, 700, 702, 706, 710,
725, 729, 731
Farge, Yves 501, 502
Fe2O3 497, 535, 541, 551, 558, 566, 572,
573, 575-577, 579, 580, 581, 583, 589,
591, 592, 596, 604, 606-608, 612, 695,
703, 717, 719, 729, 731
Fe3O4 560, 566, 572, 573, 583, 714, 731
Fendorf, Scott 497, 628, 637, 689, 701, 707,
710, 712, 715, 723
Fenter, Paul 528, 554, 555, 557-559, 575,
652, 701, 703, 707, 715, 723, 724, 727,
729, 731, 732
ferrihydrite 484, 499, 591, 592, 611,
613-619, 624-626, 670, 671, 697,
702-704, 706, 708-710, 713, 717-719,
721, 723, 724, 729, 730, 734
Ferrihydrite 644
Fourier transform infrared
spectroscopy 534
Fritsch, Emmanuel 506, 623, 656, 696, 697,
706, 708
FTIR 547, 596, 598-601, 630, 631, 685, 696,
710, 712, 716, 731, 733
GEOCHEMICAL PERSPECTIVES | V O L U M E 1 , N U M B E R 4 A N D 5
G
galena 586, 629, 634
Galoisy, Laurence 492, 502, 503, 507, 677,
678, 702, 707, 708, 716, 723, 730
Gaussian pair-distribution function 532
Gehlen, Marion 663-666, 709, 715
Gibbs free energy of adsorption 550
Gibbs, Jerry 490, 509, 510, 514-517, 519,
540, 568, 577, 683, 700, 702, 709, 719
glass alteration 677, 680, 681
glass structure 501, 503, 676, 677, 727
goethite 533, 536-539, 558, 582, 583, 591,
592, 623, 625, 629-631, 670, 696, 697,
699, 704, 709, 711-713, 715, 721, 722,
727, 731
Goldschmidt, Victor Moritz VII, 483,
485-487, 509-511, 515, 519, 521, 523,
683, 709, 710
Gouy-Chapman theory 551
grazing incidence XAFS 605, 697, 724
gypsum 488, 489, 708, 728, 729
H
haematite 511, 537, 540, 541, 558, 572,
573, 576, 577, 581, 710, 713, 714, 717,
718, 721, 722, 726-730
Harlow, George 491
Hartree-Fock 516
Hartree-Fock-level quantum mechanical
calculations 517
Helmholtz-Guoy-Chapman-SternGrahame 542, 552
Hochella, Mike IX, 488, 491, 493, 494, 499,
524, 573, 611, 612, 690, 711-714, 717
hydrolysis 543, 583, 586, 651, 653, 654,
658, 665, 677, 689, 704, 718
hydrous aluminum oxide (HAO) 581
hydrous ferric oxide (HFO) 581
hydroxylation 546, 566, 569, 572-574, 659
I
Ildefonse, Philippe 505-507, 626, 632, 665,
678, 695, 702, 712, 713, 717, 719, 720,
723
immobile element 656, 660
infrared (IR) spectroscopy 546
inner Helmholtz plane (IHP) 550
inner-sphere complexes 536, 537, 542,
550, 553, 555, 556, 592, 595, 624, 626,
631, 644, 647
J
Juillot, Farid 497, 506, 507, 639, 670, 671,
702, 706, 713, 719, 720, 721, 729
K
kaolinite 505, 506, 512, 534, 547, 582,
625-627, 658, 665, 696, 718, 720, 721,
725
K-edge 529-531, 533, 534, 536, 537, 563,
594, 615-617, 619, 641, 643, 645, 664,
665, 668, 670, 679, 702, 704, 716, 730
Krauskopf, Konrad VII, 483, 486, 487, 491,
509, 520,-522, 524, 663, 667, 669, 683,
684, 715
L
Langmuir adsorption isotherm 520
Langmuir, Irving VII, 483, 485-487, 494,
520, 524, 569, 626, 683, 684, 697, 699,
706, 710, 712, 713, 715-717, 720, 724,
725, 727-729, 731, 732, 735
laterite 505, 506, 656, 667, 669, 708, 720
leached layer 651-653, 674, 686, 711
Lead 586-588, 591, 628, 632-634, 721
Leadville 628, 629, 631, 632, 722
Leblanc, Marc 644, 647, 706, 716, 720
Levard, Clement 499, 554, 619-622, 704,
712, 714, 716, 730
Linac Coherent Light Source (LCLS) 527,
627, 735
M
macroscopic uptake measurements 581,
587
magnetite 544, 558, 585, 592-594, 612,
714, 720, 721, 729
Maher, Kate 499, 650, 652, 662, 673, 717
Manceau, Alain 503, 535, 538, 587, 591,
613, 615, 616, 628, 667, 703, 704, 706,
718
marcasite 616
mass-transfer coefficient 656
GEOCHEMICAL PERSPECTIVES | G O R D O N E . B R O W N J R . G E O R G E S C A L A S
739
N
nanoparticles 484, 488, 498, 499, 539, 558,
585, 592, 594, 596, 611-615, 619, 620,
622, 623, 627, 685, 688, 690, 692, 696,
697, 699, 703, 709, 710, 712, 716, 717,
719, 726, 729, 730
near edge X-ray absorption fine structure
(NEXAFS) 562
nickel 693, 706, 708, 714, 718
Nilsson-Pettersson model 564, 565
NMR spectroscopy 617, 718
NOM 509, 596, 599, 602, 603, 607, 609,
610, 616, 618, 623-625, 633-635, 641,
649, 685, 686, 689, 693, 735
Noyelles-Godault IV, 632
Np(V) 493, 581, 582
nuclear fuel cycle 675, 706
nuclear glass 678, 703, 708
nuclear glasses 507, 508, 675-677, 682, 702
740
O
O1s XPS 568, 569, 570, 571, 573
Ona-Nguema, Georges 498, 507, 545, 592,
649, 720, 721, 729
Ostergren, John 492, 534, 535, 591,
628-632, 701, 719, 722, 727
outer Helmholtz plane (OHP) 550, 552,
735
outer-sphere complexes 505, 534, 536,
538, 552, 553, 603, 626
oxalate 596-602, 641, 696, 713, 731
oxygen K-edge NEXAFS 562, 563
P
Pahokee peat humic acid (PPHA) 599
pair distribution function (PDF) 613
Parise, John 526, 613, 719, 722
Parks, George 487, 492, 493, 524, 533, 546,
547, 553, 558, 567, 591, 629, 696, 697,
700, 701, 703, 707, 708, 711, 713, 717,
721, 722, 724, 727, 728
Pauling bond-valence 541, 565, 588
Pauling, Linus VII, 483, 486, 487, 490,
509, 510, 511, 512, 513-517, 519, 656,
683, 713, 722
Paulings rules 502, 507, 511-513, 659, 676,
677, 681
Paulings Rules 487
Pb sorption 631, 724
PDF analysis 614, 619, 718
photoelectron multiple scattering 626
pH point of zero charge (pH PZC) 536, 550
plumbogummite 586, 634, 635
polyacrylic acid (PPA) 606, 607, 735
polywater 561, 562, 696, 697, 705, 717, 724
potentiometric titrations 542
pressure gap 687, 688
Prewitt, Charlie 491, 509, 510, 516, 517,
709, 725
Putnis, Andrew and Christine 652, 683
Q
quantum chemical modelling 566, 588,
653
quantum mechanics 526, 540, 683, 722,
728
quantum theory 540, 696
quartz 520, 546, 547, 558, 559, 566, 582,
653, 654, 696, 706, 708, 715, 719, 722,
724, 725, 731
GEOCHEMICAL PERSPECTIVES | V O L U M E 1 , N U M B E R 4 A N D 5
R
radial distribution function 531, 537, 538
radiation-damage 660, 705
radionuclides 505, 626, 627
radius ratio 487, 509, 510, 513
reductive dissolution 651, 661
reprecipitation 621, 651, 652, 674, 686,
690, 711
resonant anomalous X-ray reflectivity
(RAXR) 535
Rietveld refinement 643
Rio Negro 506, 623, 624
Rntgen, Wilhelm Conrad 526, 562, 724
rutile 540, 542, 543, 554, 556, 558, 659,
687, 697, 698, 707, 708, 713, 717, 731,
732
S
schwertmannite 645, 646, 647, 702, 706
seawater 487, 509, 521, 522, 551, 583, 663,
665, 671, 684, 714, 716, 720
selenate536-538,723
selenite 536, 537, 538, 610, 701
selenium 538, 610, 711, 717, 723, 727
Shewanella oneidensis 606-608, 638, 710,
729
silver nanoparticles 484, 499, 611, 619,
712, 716
SIMS 674, 736
size gap 687
SLAC National Accelerator
Laboratory VIII, 491, 527, 573, 627, 735
Smith, Joe 495, 504, 516, 704, 714, 725,
726
soils 485, 503-506, 535, 539, 545, 587, 602,
603, 625, 628, 631-634, 636, 638, 639,
641-644, 650-652, 656, 657, 660, 663,
667-670, 683, 685, 688, 693, 696, 701,
703-706, 713-715, 718-720, 723, 726,
727, 730, 732
sorption complexes 529, 536, 552, 554,
584, 631, 641, 643, 669
Stanford Synchrotron Radiation
Lightsource (SSRL) VIII, 491, 736
Stumm, Werner VII, 483, 485-487, 493,
520, 523-525, 536, 551-553, 566, 581,
583, 584, 587, 595, 598, 601-603, 650,
651, 656, 661, 683-685, 690, 699, 702,
708, 712, 715, 719, 725, 726, 731, 732
STXM 495, 616, 617, 674, 698, 736
T
temperature programmed desorption
(TPD) 547
Templeton, Alexis 496, 528, 545, 554, 603,
604, 606-610, 653, 727, 728
time-of-flight neutron diffraction 557
Tingle, Tracy 493, 629, 708, 722, 724
TiO2 542, 543, 551, 560, 572, 583, 612, 687,
697, 705, 713-715, 717, 723, 731, 732
tooeleite 644, 645, 646, 647, 706, 720
Trainor, Tom IX, 488, 492, 496-498, 528,
540, 554, 575, 577-580, 589, 591, 604,
697, 706, 709, 710, 717-719, 722, 727,
728, 730
transmission electron microscopy 534,
592, 698, 734, 736
Triple Layer model 552, 736
U
uranium 501, 506, 625, 626, 671, 675, 695,
696, 700, 702, 716, 729
V
van der Waals interactions 564, 719
vermiculite 557
W
water adsorption 573, 659, 696, 721, 731
water structure models 562, 731
Waychunas, Glenn 492, 503, 545, 611, 619,
637, 641, 647, 700, 702, 706, 709, 713,
721, 722, 724, 726-730, 732
weathering 484, 485, 487, 503, 517, 522,
582, 601, 625-627, 634, 642, 650, 652,
653, 656, 657, 660-662, 667, 669, 675,
683, 699, 708, 711, 715, 717, 720, 721,
723-726, 729, 731, 732
GEOCHEMICAL PERSPECTIVES | G O R D O N E . B R O W N J R . G E O R G E S C A L A S
741
X
XAFS spectroscopy 492, 495, 502, 507,
529, 531-536, 538, 539, 553, 585,
587-589, 592, 593, 607, 619, 622, 623,
629, 637, 670, 671, 685
Xanes 721
XANES 502, 503, 530, 531, 592, 594, 609,
615-617, 641, 643, 645, 664, 666, 668,
678-680, 700, 701, 705, 706, 720, 726,
731, 734, 736
XPS 493-495, 566-575, 588, 604, 619, 629,
657, 667, 674, 685, 720, 736
X-ray diffraction 510-512, 514, 526, 527,
613, 614, 676, 718, 724
X-ray Diffraction 736
X-ray fluorescence 528, 529, 534, 604, 622,
697, 714
X-ray Fluorescence 736
742
Z
zeta potential 542, 620
zinc 488, 632, 711, 716, 717, 720, 721
zircon 484, 506, 656-661, 680, 696, 705,
706, 707, 724
zirconium 656, 658, 680, 707, 708
ZnS 499, 612, 714
GEOCHEMICAL PERSPECTIVES | V O L U M E 1 , N U M B E R 4 A N D 5
GEOCHEMICAL PERSPECTIVES | G O R D O N E . B R O W N J R . G E O R G E S C A L A S
743
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ISSN 2223-7755 (print)
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DOI 10.7185/geochempersp.1.4
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University of Copenhagen,
Denmark
Marie-Aude Hulshoff
Graphical Advisor
Juan Diego
RodriguezBlanco
University of Leeds, UK
President
Vice-President
Past-President
Treasurer
Secretary
Goldschmidt Officer
Goldschmidt Officer
www.eag.eu.com
Mineral-Aqueous
Solution Interfaces
and Their Impact
on the Environment