Download 4550-15Lecture23 - Earth and Atmospheric Sciences

Mineral Surface
Reactions(cont.);
Trace Element
Geochemistry
Lecture 23
Review: Surfaces in Water
• Oxygen and metal atoms at an
oxide surface are incompletely
coordinated hence have partial
charge.
• Consequently, mineral surfaces
immersed in water attract and bind
water molecules. The water
molecules then dissociate, leaving
a hydroxyl group bound to the
surface metal ions:
≡M+ + H2O ⇄ ≡MOH + H+
• Similarly, unbound oxygens react
with water to leave a surface
hydroxyl group:
≡O– + H2O ⇄ ≡OH + OH–
• The surface quickly becomes
covered with hydroxyls (≡SOH),
considered part of the surface
rather than the solution.
Adsorption of Metals and
Ligands
• Solutes in water can
then be adsorbed to
the surface as well.
Development of Surface
Charge
• Mineral surfaces develop
electrical charge for three
reasons:
o Complexation reactions between
the surface and dissolved species,
such as those we just discussed.
Most important among these are
protonation and deprotonation.
This aspect of surface charge is pHdependent.
o Lattice imperfections at the solid
surface as well as substitutions
within the crystal lattice (e.g., Al3+
for Si4+). Because the ions in
interlayer sites of clays are readily
exchangeable, this mechanism is
particularly important in the
development of surface charge in
clays.
o Hydrophobic adsorption, primarily
of organic compounds, and
“surfactants” in particular
(discussed in Chapter 12).
Surface Charge
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We define σnet as the net density of electric charge on the solid
surface, and express it as:
σnet = σ0 + σH + σSC
where σ0 is the intrinsic surface charge due to lattice
imperfections, etc., σH is the net proton charge (i.e., the net of
binding H+ and OH–), σSC is the charge due to other surface
complexes. σ is usually measured in coulombs per square meter
(C/m2). σH is given by:
σH = F(ΓH - ΓOH)
where F is the Faraday constant and ΓH and ΓOH are the
adsorption densities (mol/m2) of H+ and OH– respectively. In a
similar way, the charge due to other surface complexes is given
by
σSC = F(zMΓM + zAΓA)
where the subscripts M and A refer to metals and anions
respectively, and z is the charge of the ion.
Thus net charge on the mineral surface is:
σnet= σ0+ F(ΓH–ΓOH+ zMΓM+ zAΓA)
Charge as a function of pH
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At some value of pH the surface
charge, σnet, will be zero. The pH at
which this occurs is known as the
isoelectric point, or zero point of
charge (ZPC). The ZPC is the pH at
which the charge on the surface of
the solid caused by binding of all
ions is 0, which occurs when the
charge due to adsorption of
cations is balanced by charge due
to adsorption of anions.
A related concept is the point of
zero net proton charge (pznpc),
which is the point of zero charge
when the charge due to the
binding of H+ and OH– is 0; that is,
pH where σH = 0.
Surface charge depends upon the
nature of the surface, the nature of
the solution, and the ionic strength
of the latter.
o
ZPC, however does not depend on ionic
strength.
Determining Surface Charge
• The surface charge due to binding of protons and
hydroxyls is readily determined by titrating a solution
containing a suspension of the material of interest
with strong acid or base. The idea is that any deficit
in H+ or OH– in the solution is due to binding with the
surface.
Surface Potential
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The charge on a surface exerts a force on ions in the adjacent solution
and gives rise to an electric potential, Ψ (measured in volts), which in turn
depends upon the nature and distribution of ions in solution, as well as
intervening water molecules.
The surface charge, σ, and potential at the surface, Ψ0, can be related
by Gouy-Chapman theory, which is similar to Debye–Hückel theory. The
relationship between surface charge and the electric potential is:
e2 x -1
sinh x =
where z is the valence of a symmetrical
2e x background electrolyte (e.g., 1
for NaCl), Ψ0 is the potential at the surface, F is the Faraday constant, T is
temperature, R is the gas constant, I is ionic strength of the solution in
contact with the surface, εr is the dielectric constant of water, and ε0 is
the permittivity of a vacuum. Most terms are constants, so at constant
temperature, this reduces to:
s = a I 1/2 sinh ( b zY0 )
where α and β are constants with values of 0.1174 and 19.5, respectively,
at 25˚C.
Surface Potential
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Where the potential is small,
the potential drops off with
distance from the surface as:
Y(x) = Y 0 e-k x
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where κ has units of inverse
length and is called the
Debye length:
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The inverse of κ is the
distance at which the
electrostatic potential will
decrease by 1/e.
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(Atom diameters are
~10-1 nm)
The one variable, other than
temperature, is I.
We also see the potential will drop off
more rapidly at high ionic strength.
Development of the ‘Double
Layer’
• The surface charge
results in an excess
concentration of
oppositely charged ions
Na+ in this case, and a
deficit of like charged
ions, Cl- in this case, in
the immediately
adjacent solution.
• Thus an electric double
layer develops
adjacent to the mineral
surface.
The Double Layer
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The inner layer, or Stern Layer, consists
of charges fixed the the surface.
The outer diffuse layer, or Gouy Layer,
consists of dissolved ions that retain
some freedom of thermal movement.
The Stern Layer is sometimes further
subdivided into an inner layer of
specifically adsorbed ions (inner
sphere complexes) and an outer layer
of ions that retain their solvation shell
(outer sphere complexes), called the
inner and outer Helmholtz planes
respectively.
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Hydrogens adsorbed to the surface are generally
considered to be part of the solid rather than the
Stern Layer.
The thickness of the Gouy (outer)
Layer is considered to be the Debye
length, 1/κ and will vary inversely with
the square root of ionic strength.
Thus the Gouy Layer will collapse in
high ionic strength solutions and
expand in low ionic strength ones.
Importance of the Double
Layer
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When clays are strongly compacted, the Gouy layers of
individual particles overlap and ions are virtually excluded from
the pore space. This results in retardation of diffusion of ions, but
not of water. As a result, clays can act as semi-permeable
membranes. Because some ions will diffuse more easily than
others, a chemical fractionation of the diffusing fluid can result.
At low ionic strength, the charged layer surrounding a particle
can be strong enough to repel similar particles with their
associated Gouy layers. This will prevent particles from
approaching closely and hence prevent coagulation. Instead,
the particles form a relatively stable colloidal suspension.
As the ionic strength of the solution increases, the Gouy layer is
compressed and the repulsion between particles decreases. This
allows particles to approach closely enough that they are bound
together by attractive van der Waals forces between them.
When this happens, they form larger aggregates and settle out of
the solution.
For this reason, clay particles suspended in river water will
flocculate and settle out when river water mixes with seawater in
an estuary.
Effect of the surface potential
on adsorption
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The electrostatic forces also affect complexation reactions at the
surface, as we noted at the beginning of this section. An ion must
overcome the electrostatic forces associated with the electric
double layer before it can participate in surface reactions. We
can account for this effect by including it in the Gibbs free
energy of reaction:
∆Gads=∆Gintr+∆Gcoul
where ∆Gads is the total free energy of the adsorption reaction,
∆Gint is the intrinsic free energy of the reaction (i.e., the value the
reaction would have in the absence of electrostatic forces; e.g.,
the same reaction taking place in solution), and ∆Gcoul is the free
energy due to the electrostatic forces and is given by:
∆Gcoul = F∆ZΨ0
where ∆z is the change in molar charge of the surface species
due to the adsorption reaction. For example, in the reaction:
≡SOH+Pb2+ ⇄ ≡SOPb+ +H+
the value of ∆z is +1 and ∆Gcoul =FΨ0
Effect on Equilibrium
• Thus if we can calculate ∆Gcoul, this term can be
added to the intrinsic ∆G for the adsorption
reaction (∆Gintr) to obtain the effective value of ∆G
(∆Gads). From ∆Gads it is a simple and straightforward
matter to calculate Kads:
K = e-∆ Gads /RT = e-∆ Gintrin /RT e-∆ Gcoul /RT
o (note equation 6.128 in the book has typos. Should be the above.)
• Since Kintr = e-∆Gintr/RT and ∆Gcoul = F∆zΨ0, we have:
K = K intr e- F∆ zY0 /RT
• Thus we need only find the value of Ψ0, which we
can calculate from σ.
Effect of potential on
adsorption
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The effect of surface
potential on a given
adsorbate will be to shift the
adsorption curves to higher
pH for cations and to lower
pH for anions.
The figure illustrates the
example of adsorption of Pb
on hydrous ferric oxide. When
surface potential is
considered, adsorption of a
given fraction of Pb occurs at
roughly 1 pH unit higher than
in the case where surface
potential is not considered.
In addition, the adsorption
curves become steeper.
Introduction to Trace
Element Geochemistry
Chapter 7
Importance
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Though trace elements, by definition, constitute only a small
fraction of a system of interest, they provide geochemical and
geological information out of proportion to their abundance.
There are several reasons for this.
First, variations in the concentrations of many trace elements are
much larger than variations in the concentrations of major
components, often by many orders of magnitude.
Second, in any system there are far more trace elements than
major elements. In most geochemical systems, there are ten or
fewer major components that together account for 99% or more
of the system. This leaves 80 trace elements. Each element has
chemical properties that are to some degree unique, hence
there is unique geochemical information contained in the
variation of concentration for each element. Thus the 80 trace
elements always contain information not available from the
variations in the concentrations of major elements.
Third, the range in behavior of trace elements is large, and
collectively they are sensitive to processes to which major
elements are insensitive.
What is a Trace Element?
• For most silicate rocks, O, Si, Al, Na, Mg, Ca, and Fe are “major
elements”.
• An operational definition might be as follows: trace elements
are those elements that are not stoichiometric constituents of
phases in the system of interest.
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This definition is a bit fuzzy: a trace element in one system is not one in another. For
example, K never forms its own phase in mid-ocean ridge basalts (MORB) but K is
certainly not a trace element in granites.
Of course this definition breaks down for fluid solutions such as water and magma.
• Trace elements in seawater and in rocks do have one thing in
common: neither affect the chemical or physical properties of
the system as a whole to a significant extent.
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But some components do influence properties even at very low conc.; e.g. trace
elements can influence color of minerals.
• Perhaps the best definition of a trace element is: an element
whose activity obeys Henry’s Law in the system of interest.
o
This implies sufficiently dilute concentrations that, for trace element A and major
component B, A–A interactions are not significant compared to A–B interactions,
simply because A–A interactions will be rare.
Concentrations in the Silicate
Earth
Goldschmidt’s Classification
• Atmophile elements are generally extremely volatile (i.e., they
form gases or liquids at the surface of the Earth) and are
concentrated in the atmosphere and hydrosphere.
• Lithophile, siderophile and chalcophile refer to the tendency
of the element to partition into a silicate, metal, or sulfide
liquid respectively.
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Goldschmidt classified them based on the minerals in which they were
concentrated in meteorites.
• Lithophile elements are those showing an affinity for silicate
phases and are concentrated in the silicate portion (crust and
mantle) of the Earth.
• Siderophile elements have an affinity for a metallic liquid
phase. Concentrated in Earth’s core.
• Chalcophile elements have an affinity for a sulfide liquid
phase. They are also depleted in the silicate earth and may
be concentrated in the core.
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Most elements that are siderophile are usually also somewhat chalcophile and vice
versa.
Goldschmidt’s Classification
Distribution of the Elements
• Atmophile in the
atmosphere.
• Siderophile (and
perhaps chalcophile) in
the core.
• Lithophile in the mantle
and crust.
o Incompatible elements, those
not accepted in mantle
minerals, are concentrated in
the crust.
o Compatible elements
concentration in the mantle.
Geochemical Classification
Geochemical Classification
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The volatile elements: Noble gases, nitrogen
The semi-volatiles
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The alkali and alkaline earth elements
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So called because of their high ionic charge: Zr and Hf have +4 valence states and Ta and Nb have +5 valence states.
Th (+4) and U (+4, +6) are sometimes included in this group. Because of their high charge, all are relatively small cations
but because of their high charge they are excluded from most minerals and hence are also incompatible.
The first series transition metals
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The alkali and alkaline earth elements have a single valence state (+1 for the alkalis, +2 for the alkaline earths). The
bonds these elements form are strongly ionic in character (Be is an exception, as it forms bonds with a more covalent
character)These elements relatively soluble in aqueous solution. The atoms of these elements behave approximately
as hard charged spheres; size and charge dictate their behavior in igneous processes.
The heavier alkalis and alkaline earths have ionic radii too large to fit in many minerals (mantle ones in particular). As a
consequence, they are said to be incompatible.
High field strength (HFS) elements
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they partition readily into a fluid or gas phase (e.g., Cl, Br) or form compounds that are volatile (e.g., SO2, CO2). Not all
are volatile in a strict sense (volatile in a strict sense means having a high vapor pressure or low boiling point; indeed,
carbon is highly refractory in the elemental form).
Chemistry is considerably more complex: many of the transition elements have two or more valence states; covalent
bonding is more important (bonding with oxygen in oxides and silicates is predominantly ionic, but bonding with other
non-metals, such as sulfur, can be largely covalent) A final complicating factor is the geometry of the d-orbitals, which
are highly directional and thus bestow upon the transition metals specific preferences for the geometry of
coordinating ligands. They range from moderately incompatible (e.g., Ti, Cu, Zn) to very compatible (e.g., Cr, Ni, Co) .
The noble metals
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The platinum group elements (Rh, Ru, Pd, Os, Ir, Pt) plus gold (and Re) are often collectively called the noble metals.
These metals are so called for two reasons: first, they are rare, and second, they are unreactive and quite stable in
metallic form. Their rarity is in part a consequence of their highly siderophilic character. These elements are all also to
varying degrees. These elements exist in multiple valence states, ranging from 0 to +8, and have bonding behavior
influenced by the d-orbitals.
Au Pt, Re, and Pd are moderately incompatible elements, while Rh. Ru, Os, and Ir are highly compatible
The Rare Earth Elements
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The rare earths and Y are strongly
electropositive. As a result, they form
predominantly ionic bonds, and behave as
hard charged spheres.
The lanthanide rare earths are in the +3
valence state over a wide range of
oxygen fugacities.
In the transition metals, the s orbital of the
outermost shell is filled before filling of lower
electron shells is complete so the
configuration of the valence electrons is
similar in all the rare earth, hence all exhibit
similar chemical behavior.
Ionic radius, which decreases progressively
from La3to Lu3+ (93 pm) governs their
relative behavior.
Because of their high charge and large
radii, the rare earths are incompatible
elements.
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The degree of incompatibility of the lanthanides varies
with atomic number. Highly charged U and Th are highly
incompatible elements, as are the lightest rare earths.
However, the heavy rare earths have sufficiently small
radii that they can be accommodated to some degree in
many common minerals such as Lu for Al in garnet. Eu2+
can substitute for Ca in plagioclase.
Rare Earth Diagrams
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The systematic variation in lathanide rare
earth behavior is best illustrated by plotting
the log of the relative abundances as a
function of atomic. Relative abundances
are calculated by dividing the
concentration of each rare earth by its
concentration in a set of normalizing values,
such as the concentrations of rare earths in
chondritic meteorites.
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Rare earths are also refractory elements, so that their
relative abundances are the same in most primitive
meteorites - and presumably (to a first approximation) in
the Earth.
Why do we use relative abundances? The
abundances of even-numbered elements
in the solar system are greater than those of
neighboring odd-numbered elements and
abundances generally decrease with
increasing atomic number, leading to a
saw-toothed abundance pattern.
Normalizing eliminates this.
Abundances in chondritic meteorites are
generally used for normalization. However,
other normalizations are possible: sediments
(and waters) are often normalized to
average shale.