Solubility (cont.);
Mineral Surfaces
& Reactions
Lecture 22
Constructing stability
diagrams
ΣCO2 = 5x 10-2 M
• This diagram shows the stability
of ferrous iron minerals as a
function of pH and sulfide for
fixed total Fe and CO2.
• Procedure: manipulate
equilibrium constant
expressions to obtain and
expression for ΣS in terms of pH.
For example:
[Fe2+] = 10-6 M
(Pyrrhotite)
(Siderite)
o FeCO3 + H+ ⇋ Fe2+ + HCO3–
o FeCO3 + 2H2O ⇋ HCO3- H+ + Fe(OH)2
o FeS + 2H2O ⇋ Fe(OH)2 + H+ + HS-
• Trick: simplify by ignoring
pH = log K FeCO - log[HCO3- ]- log[Fe2+ ] = 7.5
②
species present at low conc.
(e.g., CO32- at low pH).
③ pH = log[HCO3 ]+13.0
3
- pH
➄ log SS = pH - pK FeS + pK Fe(OH )2 + log(K S +10 )
Solubility of SiO2
• Silica forms silicic acid
(H4SiO4) in solution, which
can then dissociate through
a series of reactions, e.g.,
• H4SiO4 ⇋ H3SiO4– + H+
K1
• H3SiO4– ⇋ H2SiO42- + H+ K2
• Solubility can be expressed
as:
ìï K1 K1K 2 üï
[SiO2 ]T = [H 2 SiO4 ] í1+ 2 ý
a
aH + þï
+
H
îï
• where [H4SiO4] is controlled
by solubility of either quartz
or amorphous silica.
• As a consequence, its
solubility is a function of pH:
high only at high pH.
Solubility of Hydroxides
•
•
The hydroxide is the least soluble salt of many metals. Therefore,
it is the solubility of their hydroxides that controls the solubility of
these metals in natural waters.
Since these dissolution reactions involve OH–, they are pHdependent, and the slope of the solubility curve depends on
the valence of the metal (e.g., -3 for Fe3+, -2 for Fe2+, -1 for Ag+).
Solubility of Al(OH)3
• Solubility of gibbsite:
o
Al(OH)3 + 3H+ ⇋ Al3+ + H2O
K gib =
aAl 3+
aH +
= 10 -8.1
• However, Al forms hydroxide
complexes, e.g.:
• Al3+ + H2O ⇋ Al(OH)2+ + H+
• The total dissolved Al will be
the sum of all Al species in
solution:
ìï K
K
K
K üï
aAl 3+T = aAl 3+ í1+ 1 + 2 2 + 3 3 + 4 4 ý
ïî aH + aH + aH + aH + ïþ
• A consequence of this is that
acid rain leads to Al
poisoning.
Solubility of Ferric Iron
Silicate Solubility
• We’ve looked at the solubility of Si, Al, Fe and other
cations in the isolation of simple laboratory-like
systems.
• The real world is usually more complex. Silicate rocks
predominate at the surface of the Earth, thus Si, Al,
and other cations will generally all be present.
• In addition to gibbsite and SiO2, some of the more
common weathering products of silicate rocks
include:
• kaolinite: Al2Si2O5(OH)4
• pyrophyllite: Al2Si4O10(OH)2
• illite (muscovite): KAl3Si3O10(OH)2
Silicate Solubility
Gibbsite will precipitate from Al-bearing solution only at lowest
concentrations of SiO2. Occurrence generally restricted to highly
weathered soils where all the SiO2 has washed out.
Silicate Solubility
Clay Minerals
• Review section 6.5 to become familiar with clay
minerals, but we will not cover it in class.
Adsorption and Surface
Complexation
• We can define adsorption as attachment of an ion in solution
to a pre-existing solid surface. It involves one or more of the
following:
• Surface complex formation: The formation of coordinative
bonds between metals and ligands at the surface, similar to
the formation of complexes in solution.
• 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.
• Hydrophobic adsorption: Many organic substances, most
notably lipids, are highly insoluble in water due to their nonpolar nature. These substances become adsorbed to surfaces,
not because they are attracted to the surface, but rather
because they are repelled by water.
o
This topic is treated in Chapter 12.
Surface Complexation Model
QM =
K ad [M ]
1+ K ad [M ]
•
The Langmuir Isotherm was
•
and the Freundlich was:
•
The tendency of an ion to be adsorbed to a surface is expressed as an
adsorption coefficient, which we can relate to thermodynamics and
energy and entropy changes at the mineral surface, i.e., ∆Gad.
The surface complexation model incorporates both chemical bonding of
solute species to surface atoms and electrostatic interactions between
the surface and solute ions.
The free energy of adsorption is the sum of a complexation, or intrinsic,
term and an electrostatic, or coulombic term:
∆Gad = ∆Gintr + ∆Gcoul
From this it follows that the adsorption equilibrium constant can be written
as:
Kad = Kintr Kcoul
•
•
•
QM = K ad [M ]
Surfaces in Water
•
•
•
•
•
•
Consider a simple surface such as a metal
oxide.
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. These hydroxyls
can then act as either proton acceptors or
proton donors through further association or
dissociation reactions:
≡SOH + H+ ⇄ ≡SOH2+
≡SOH ⇄ SO- + H+
We should not be surprised to find that these
kinds of reactions are strongly pH-dependent.
Adsorption Mechanisms
• Adsorption of metals to the
surface may occur through
replacement of a surface
proton.
• Ligands may be absorbed
by replacement of a
surface OH group.
• The adsorbed metal may
bind an additional ligand.
• The adsorbed ligand may
bind an additional metal.
• An additional possibility is
multidentate adsorption,
where a metal or ligand is
bound to more than one
surface site.
Multidentate Adsorption
• raises an interesting dilemma for the Langmuir
isotherm. Where x sites are involved, we could write
the reaction:
x≡S + M ⇄ ≡SxM
• Writing an equilibrium constant expression for this
reaction would imply that the probability of finding
x sites together is proportional to the xth power of
concentration, which is not the case.
• A better approach is to assume that the reaction
occurs with a multidentate surface species, ≡Sx and
that its concentration is [≡S]/x. The equilibrium
constant is then:
K ad =
[º Sx M ]
[M ][º S] / x
Cation pH dependence
•
•
•
•
Adsorption of metals and ligands will
be strongly pH-dependent. Adsorption
of cations increases with increasing pH.
The figure shows that adsorption of
metals on goethite goes from
insignificant to nearly complete over a
very narrow range of pH.
This reflects protonation of the surface,
but it also reflects the extent of
hydrolysis of the ion in solution.
Metals vary greatly in how readily they
are adsorbed. At a pH of 7, for
example, and a solution containing a
1 µM concentration of the metal of
interest, the fraction of surface sites
occupied by Ca, Ag, and Mg is trivial
and only 10% of surface sites would be
occupied by Cd. At this same pH,
however, 97% of sites would be
occupied by Pb and essentially all sites
would be occupied by Hg and Pd.
Anion pH dependence
• Adsorption of
anions decreases
with increasing pH.
• Extend of
adsorption also
depends on the
nature of the
ligand.
Inner & Outer Sphere
Complexes
•
•
•
•
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 (which we’ll get to shortly)
by long-range electrostatic
forces.