Download Aquatic Geochemistry

The Geochemistry of Rocks and
Natural Waters
Course no. 210301
Aquatic Geochemistry
A. Koschinsky
Aquatic Geochemistry - Introduction
Some general definitions
Concentration and activity
Molarity (mmol/l) = concentration (mg/l) / atomic weight
Molality = Molarity / density of solution (or seawater, respectively)
Excercise: 1. Express „5 µmol/l Pb“ in the unit mg/l (atomic weight Pb = 207).
2. Express „60 mg/l CO2“ in the unit mmol/l (molecular weight CO2 = 44).
Activity a = c(concentration) x f
(activity coefficient, also called )
The activity can be regarded as
effective concentration (i.e. the
share of ions in solution that is
able to react).
This form of the extended Debye-Huckel equation is valid
for 25°C and solution ionic strength up to 0.1 molar
Aquatic Geochemistry - Introduction
Activity coefficients are a function of the following:
a. hydrated radius of
ion (a)
b. charge on ion (z)
c. ionic strength of
solution (m)
m= 1/2 cizi2
Aquatic Geochemistry - Introduction
Speciation
Chemical Speciation describes the chemical form in which a compound is
present in the system; dissolved species include hydrated free cations and
different complex forms.
Importance of Speciation
A. Toxicity and bioavailability of Metals is related to the form of an ion in
solution (e.g., Cu2+ is toxic while CuCO30 is not).
B. Productivity of waters is related to the organic ligand concentration (e.g.,
organic conditioning in upwelling waters).
C. Solubility of metals is related to the form in solution (e.g., ligands can
solubilize metals from sediments).
Aquatic Geochemistry - Introduction
Speciation
Factors Controlling the State of Metal Ions
1. Oxidation State
2. pH
3. Composition of Inorganic Ligands (OH-, CO32-, HS-)
4. Composition of Organic Ligands (Humic and Fulvic Acids, others)
5. Pressure and Temperature
Importance of pH in Speciation
HX  H+ + XX- is a ligand that can complex a metal (CO32-, OH-, HPO42-, PO43-,
organic acid anions)
Aquatic Geochemistry - Introduction
•Most reactions in aqueous solutions can be placed in one of these categories:
• Acid-base reactions, e.g., dissociation of carbonic acid:
• H2CO3  H+ + HCO3
• Give the dissociation reaction for another acid:
• Complexation, e.g., hydrolysis of mercury:
• Hg2+ + H2O  Hg(OH)+ + H+
• Give the complexation reaction for dissolved Zn and chloride:
• Dissolution/Precipitation, e.g., dissolution of orthoclase:
• KAlSi3O8 + H+ + 7H2O  Al(OH)3 + K+ + 3H4SiO4
• Give the precipitation reaction of calcite:
• Adsorption/Desorption, e.g., adsorption of Mn on a clay surface:
• =S + Mn2+  =S–Mn (where =S indicates the surface of the clay).
• Give the adsorption reaction of copper on a goethite surface:
Aquatic Geochemistry - Complexation
Ions in solution often associate with other ions, forming complexes. This affects
the solubility and reactivity of ions. For example, if ions form stable, soluble
complexes, this greatly enhances the solubility of the ions.
Central ion (metal) - ligands (ions or molecules that surround, or coordinate, the
central ion)
The simplest complexes are those formed between metals and water: aquo
complexes (Water molecules are the ligands). Aquo-complexes are ubiquitous:
all charged species will have a solvation shell. Truly “free ions” do not exist.
Example: zinc aquo complex Zn(H2O)62+
Other complexes include:
• Inorganic complexes (such as chloride complexes of the form MnCl+, carbonate
complexes such as Nd(CO3)+ or hydroxide complexes such as Zr(OH)5+ and
Ti(OH)40
• Organic complexes of mostly complicated structure
Complexation - Types of complexes
Beyond aquo complexes, we
can distinguish two types of
complexes:
 Ion pairs (outer sphere
complexes), where ions of
opposite charge associate
with one and other through
electrostatic attraction, yet
each ion retains part or all of
it’s solvation sphere.
 Complexes (senso stricto,
inner sphere complexes),
where the two ions are in
contact and a bond forms
between them that is at least
partly covalent in nature.
Complexation - Stability Constants
Complex formation reaction between a metal cation M and an anion or ligand L :
mM + nL <--> MmLn
As with any other reaction, the equilibrium constant is defined as:
K = [MmLn] / [M]m [L]n
Example: Zn(H2O)62+ + OH- <---> Zn(H2O)5(OH)+ + H2O
Or:
Zn2+ +
OH- <---> Zn(OH)+
K1 = [ZnOH+] / [Zn2+] [OH]
The zinc ion might associate with a second hydroxyl:
Zn(OH)+ + OH- <---> Zn(OH)20
Define the second equilibrium constant K2 =
The equilibrium constant for complex formation is often referred to as stability
constant .
Complexation - Water-related complexes
Aquo-complexes can act as weak acids because the positive charge of the
central ion tends to repel hydrogen atoms in the water molecules:
The repulsion between the central
ion and the protons depends on:
 pH
 size of the central ion
 charge of the central ion.
Complexation - Other complexes
Complexation - Other complexes
With respect to complex formation, the elements can be divided into four classes:
1. non-metals (forming anions), ligand formers.
2. “A-type” or “hard” metals. These metals have spherically symmetric,
inert-gas type outer electron configurations. Their electron shells are not readily
deformed by electric fields and can be viewed as “hard spheres”.
They preferentially form
complexes with F and ligands
having O as the donor atoms
(OH–, CO32- , PO43- , SO42-).
Stability of the complexes formed
by these metals increases with
charge to radius ratio. Thus the
alkalis form only weak, unstable
complexes, while elements such as
Zr4+ form very strong, stable
complexes. Anions and cations are
bound primarily by electrostatic
forces, i.e., ionic-type bonds.
Complexation - Other complexes
3.
4.
B-type, or “soft”, metal ions: Their electron shells are not spherically
symmetric and are readily deformed by the electrical fields of other ions
(“soft”). They preferentially form complexes with bases having S, I, Br, Cl, or
N (such as ammonia) as the donor atom. Bonding between the metal and
ligand(s) is primarily covalent and is comparatively strong. Thus Pb form
strong complexes with Cl– and S2–.
First series transition
metals: Their electron
sheaths are not spherically
symmetric, but they are not
so readily polarizable as
the B-type metals.
On the whole, however,
their complex-forming
behavior is similar to that of
the B-type metals.
Complexation - Transition metal complexes
Among the transition metals, the
sequence of complex stability is
Mn2+ < Fe2+ < Co2+ < Ni2+< Cu2+ >
Zn2+, a sequence known as the
Irving-Williams Series. In the
figure, all the sulfate complexes
have approximately the same
stability, a reflection of the
predominately electrostatic
bonding between sulfate and
metal. Pronounced differences
are observed for organic ligands.
Although the absolute value of
stability complexes varies from
ligand to ligand, the relative
affinity of ligands having the
same donor atom for these
metals is always similar.
Complexation - Chelation
Chelation: Organic molecules can often
have more than one functional group and
hence can coordinate a metal at several
positions, a process called chelation. Such
ligands are called multi-dentate and organic
compounds having these properties are
referred to as chelators or chelating agents
(for example EDTA).
Complexation - Speciation Calculation
Complexation Speciation
Calculation
Dependence on pH and
temperature
Complexation - Speciation Calculation
Dependance on ligand concentration and temperature
Pb
0
0
Zn
1
2
4
2
4
3
3
1
Pb
1
0
2
3
Zn
1
2
4
Solid - Solution Interactions
Mineral surface - interface - solution
The processes at the interface govern equilibria between solids and solutions.
Solid - solution interactions: Adsorption
Important: Many minerals have large active surface areas (strong sorptive
properties)!
Example: In natural marine solids, the following specific surface areas were measured:
Clay fraction of a sediment:
Manganese nodule (todorokite MnO2)
Ferromanganese crust (vernadite -MnO2)
Ferrihydrite FeOOH
92 m2/g
181 m2/g
350 m2/g
140 m2/g
 Concentrations of many trace metals in many natural aqueous systems are
not controlled by precipitation/dissolution, but by adsorption/desorption on/from
mineral and organic surfaces.
 Solution complexation and adsorption are competing reactions in aquatic
systems!
Solid - solution interactions: Adsorption
Equilibria between dissolved and surface-bound species
-
Equilibration in dissolved milieu: formation of dissolved complexes by central ion and
ligands
 Stability constant  of the complex
-
Distribution of compounds between seawater and solid material (particulate
suspendended matter, bottom sediment, organisms)

Distribution or partitioning coefficient Kd
Me(diss.)<----> Me(solid)
- Competition between reaction of the ion with a ligand (solution complexation)
and reaction of the ion with a solid surface (sorption)
Adsorption mechanisms
Adsorption: attachment of an ion in solution to a preexisting solid surface (clay,
oxide, carbonate, … particle). Adsorption involves:
 Electrostatic interactions: Solid surfaces are typically electrically charged.
This electrostatic force, which is effective over greater distances than purely
chemical forces, affects surface complex formation and loosely binds other ions
to the surface.
Adsorption mechanisms
 Surface complex
formation: The formation
of coordinative bonds
between metals and
ligands at the surface
(similar to the formation
of complexes between
dissolved components).
 Hydrophobic adsorption: Many organic substances are highly insoluble in
water due to their non-polar nature. These substances become adsorbed to
surfaces because they are repelled by water.
Adsorption mechanisms
As is the case with
soluble complexes,
surface complexes may
be divided into inner
sphere and outer sphere
complexes. Inner sphere
complexes involve some
degree of covalent
bonding between the
adsorbed species and
atoms on the surface. In
an outer-sphere complex,
one or more water
molecules separate the
adsorbed ion and the
surface; in this case
adsorption involves only
electrostatic forces.
The third possibility is that an ion may be
held within the diffuse layer by long-range
electrostatic forces.
Adsorption mechanisms
The functional groups of mineral and organic surfaces have properties similar
to those of their dissolved counterparts. In this sense, surface complexation
reactions are similar to complexation reactions in solution.
Functional groups: -OH, =O (inorganic),
-COOH (carboxyl group, organic acid), …
Reactions between these surface groups and dissolved species are influenced
by the proximity of surface groups to each other. For example, the surface
charge will change systematically as the adsorbed surface concentration of a
positive species such as H+ increases. This change in surface charge will
decrease the attraction between H+ ions and the surface. As a result, the
equilibrium constant for the surface protonation reaction will change as the
surface concentration of H+ increases.
-
Oxide
-
+
H+
-->
-
Oxide
-H
-
+
H+
-->
-
Oxide
-H
-H
Equilibrium constant or distribution coefficient Kd = [Csorbed]/[Caq]
…
Outer- and inner-sphere complexes
(pK = -logK)
Surface complexation
Stable surface complexes are most likely to be formed at surface irregularities
such as kinks and steps (see scheme below, formation of copper complexes
on calcite surface). The presence of other metal ions may lead to the
incorporation of these ions into the crystal lattice, or to the inhibition of the
calcite crystal growth. Both processes involve the substitution of Ca by these
ions at the surface complexation sites.
Adsorption: oxide surfaces in water
Oxygen and metal atoms at an oxide surface are
incompletely coordinated; i.e., they are not
surrounded by oppositely charged ions as they
would be in the interior of a crystal (Fig. a).
Consequently, mineral surfaces immersed in
water bind water molecules (Fig. b), which can
then dissociate, leaving a hydroxyl group bound
to the surface metal ion:
In a similar fashion, incompletely coordinated
oxygens at the surface can also bind water
molecules, which can then dissociate, again
creating a surface hydroxyl group:
Thus the surface on an oxide immersed in water
quickly becomes covered with hydroxyl groups
(Fig. c), which are considered to constitute part
of the surface rather than the solution.
Adsorption: oxide surfaces in water
These hydroxyl groups can then act as either proton acceptors or proton donors
through further association or dissociation reactions, e.g.:
These reactions are strongly pH-dependent!
Adsorption: oxide surfaces in water
Adsorption on a mineral
surface may occur when
(a) a metal replaces a surface
proton,
or (b) a ligand replaces a
surface OH group.
The adsorbed metal (c) may
bind an additional ligand,
and the ligand (d) may bind an
additional metal.
2-
Multidentate adsorption
involves more than one
surface site (e, f).
pH dependence of adsorption
Since surface bound protons and OH groups are almost inevitably involved in
adsorption, adsorption of metals and ligands is strongly pH dependent. This strong
dependence on pH certainly reflects protonation of the surface, but it also reflects
the extent of hydrolysis of the ion in solution.
Surface charge of particles
Mineral surfaces develop electrical charge for several reasons:
 Complexation reactions between the surface and dissolved species, including
protonation and deprotonation. Because these reactions depend on pH, this aspect
of surface charge is pH dependent.
 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.
Thus there are several contributions to surface charge density. We define net as
the net density of electric charge on the solid surface, and can express it as:
net = 0 + H + SC
where 0 is the intrinsic surface charge due to lattice imperfections and
substitutions,
H is the net proton charge, i.e., the charge due to binding H+ and OH–,
SC is the charge due to other surface complexes.
 is usually measured in coulombs per square meter (C/m2).
Surface charge of particles
Surface charge of some common sedimentary materials as a function of pH. The pH
dependence of surface charge reflects the predominance of attached protons at low pH
and the predominance of attached hydroxyls at higher pH.
Surface charge of particles
At some value of pH the surface charge
will be zero. This point is known as the
isoelectric point, or zero point of charge
(ZPC). This occurs when the charge due
to adsorption of cations is balanced by
charge due to adsorption of anions.
When the pHzpc of a solid is lower than
the surrounding pH, then the solid
develops a negative surface charge.
When the pHzpc of a solid is higher than
the surrounding pH, then the solid
develops a positive surface charge.
Question: What is the surface charge of
goethite and vernadite (-MnO2) in
seawater?
2
2
Long-Term Fate of Adsorbed Ions
 Desorption of ions is much less favored than sorption.
 Sorbed ion complexes may polymerize to form surface precipitatites.
 Sorbed ions may be incorporated into the host structure via dissolution and
recrystallization.
Surface precipitation and solid solution
In surface precipitation, cations (or anions) which adsorb to the surface of a mineral,
may form a precipitate of the cation (anion) with the constituent ions of the mineral
at high surface coverage.
Cations at the solid-water interface are treated as surface species, while those not
in contact with the solution phase are treated as solid species forming a solid
solution. The formation of solid solution implies isomorphic substitution.
Surface precipitate
Solid solution
Surface precipitation and solid solution
There is a continuum of surface complexation (adsorption) and surface
precipitation. At low sorbate concentrations, surface complexation is the dominant
mechanism (Fig. a).
As the sorbate concentration increases, surface precipitation becomes the dominant
“sorption” (metal ion incorporation) mechanism (Fig. b).
Aquatic Geochemistry - Precipitation/Dissolution
A few general rules:
 Precipitation of ions as
insoluble (low solubility) minerals
is a major mechanism by which
metal concentrations are limited
in groundwater and other natural
waters.
 A useful generalization is that
highly charged cations (+2, +3,
+4) form insoluble precipitates
with highly charged anions
(sulfate, carbonate, phosphate)
 Chalcophile elements (Pb, Cd,
Cu, Zn, Hg) form insoluble
sulfides.
Dissolution/precipitation - Some definitions
Dissolution/precipitation - carbonate minerals
Dissolution/precipitation - hydroxide minerals
Dissolution/precipitation - sulfate minerals
Dissolution/precipitation - sulfide minerals
Dissolution/precipitation - solubility calculations
8
Convert the result in mol/l to ppm!
Dissolution/precipitation - solubility calculations
than PbSO4
The presence of minerals such as CaCO3 and CaSO4 can suppress the solubility
of metals such as Pb and Cd via the common ion effect.
Dissolution/precipitation - solubility calculations
2,7,9 ?
Because they are strong bases, the solubilities of hydroxides, sulfides and
carbonates will depend on the pH.