v ¼
dC
dt
¼ k  S
ð1:24Þ
where k is the rate constant (mol l
À1 m
À2 s
À1 ) and S is the surface area of the solid
(m
2 ).
One the one hand, outer sphere surface complexes show small effects on dissolution rates, while the oxidation state of the metal ion remains unchanged. On the
other hand, the surface reactivity of inner sphere complexes with ligands is enhanced
or inhibited depending on the nature of the ligand. When surface metals cations are
bound by water, hydroxide or ligands bound to one metal center, the dissolution is
easy. Alternatively, when ligands are bridging several metal centers (e.g., phosphate)
or when they are blocking the surface by hydrophobic moieties (e.g., HA, polymers),
the dissolution is inhibited. The electron density or negative charge of ligands that
enter into the metal coordination sphere decreases its Lewis acidity and enhance the
lability of the metal-surface bonds, favoring metal dissolution. The formation of
surface complexes is fast, while the release of metal ions in solution is a slow step.
There are two simultaneous ways to enhance the solubility of metals ions: (1) the free
metal ions concentration is lowered in solution since metals are bound to ligands as
soluble complexes and slow surface dissolution and interlattice reactions happen to
maintain equilibrium concentrations and (2) the adsorption of a ligand onto the solid
surface leads to the formation of ternary complexes as an activated intermediate
before the release of a soluble complex and a free site occurs at the solid surface.
According to the activated complex theory and assuming the steady-state, the
dissolution rate v increases with the surface concentration of the precursor of the
activated complex:
v ¼
dC
dt
¼ k
0
 C
Ã
ð1:25Þ
where k
0 is the rate constant (m
2 l
À1 s
À1 ) and C
à is the surface concentration of the
precursor (mol m
À2 ). Thus, there is an ambivalent role of adsorption for ligands as it
is involved in surface dissolution, but their use has to be kept low and specific
because of costs and preservation of soils properties. Therefore, the balance between
surface complexation and metal dissolution (e.g., citrate vs. EDTA) is critical when
considering ligands.
Thus, the appropriate selection of ligands and process operating conditions are
essential both in order to ensure the success of the treatment and to control side
effects. This choice builds on a large literature complicated by the wide variability of
environmental conditions. It must be supported by lab tests for matrix characterization (especially leaching tests) and treatability. Selectivity of ligands for contaminant
removal is important. It may have a kinetic or a thermodynamic origin. The kinetic
origin is related to contaminant reactivity and hence to site reactivity also if any.
When exchange kinetics are fast, selectivity is imposed by the relative stability of
contaminant–ligand complexes with respect to those with competitors and relative
concentrations. The selection of the active agents and their concentrations is built on
30
N. Fatin-Rouge
dC
dt
¼ k  S
ð1:24Þ
where k is the rate constant (mol l
À1 m
À2 s
À1 ) and S is the surface area of the solid
(m
2 ).
One the one hand, outer sphere surface complexes show small effects on dissolution rates, while the oxidation state of the metal ion remains unchanged. On the
other hand, the surface reactivity of inner sphere complexes with ligands is enhanced
or inhibited depending on the nature of the ligand. When surface metals cations are
bound by water, hydroxide or ligands bound to one metal center, the dissolution is
easy. Alternatively, when ligands are bridging several metal centers (e.g., phosphate)
or when they are blocking the surface by hydrophobic moieties (e.g., HA, polymers),
the dissolution is inhibited. The electron density or negative charge of ligands that
enter into the metal coordination sphere decreases its Lewis acidity and enhance the
lability of the metal-surface bonds, favoring metal dissolution. The formation of
surface complexes is fast, while the release of metal ions in solution is a slow step.
There are two simultaneous ways to enhance the solubility of metals ions: (1) the free
metal ions concentration is lowered in solution since metals are bound to ligands as
soluble complexes and slow surface dissolution and interlattice reactions happen to
maintain equilibrium concentrations and (2) the adsorption of a ligand onto the solid
surface leads to the formation of ternary complexes as an activated intermediate
before the release of a soluble complex and a free site occurs at the solid surface.
According to the activated complex theory and assuming the steady-state, the
dissolution rate v increases with the surface concentration of the precursor of the
activated complex:
v ¼
dC
dt
¼ k
0
 C
Ã
ð1:25Þ
where k
0 is the rate constant (m
2 l
À1 s
À1 ) and C
à is the surface concentration of the
precursor (mol m
À2 ). Thus, there is an ambivalent role of adsorption for ligands as it
is involved in surface dissolution, but their use has to be kept low and specific
because of costs and preservation of soils properties. Therefore, the balance between
surface complexation and metal dissolution (e.g., citrate vs. EDTA) is critical when
considering ligands.
Thus, the appropriate selection of ligands and process operating conditions are
essential both in order to ensure the success of the treatment and to control side
effects. This choice builds on a large literature complicated by the wide variability of
environmental conditions. It must be supported by lab tests for matrix characterization (especially leaching tests) and treatability. Selectivity of ligands for contaminant
removal is important. It may have a kinetic or a thermodynamic origin. The kinetic
origin is related to contaminant reactivity and hence to site reactivity also if any.
When exchange kinetics are fast, selectivity is imposed by the relative stability of
contaminant–ligand complexes with respect to those with competitors and relative
concentrations. The selection of the active agents and their concentrations is built on
30
N. Fatin-Rouge