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Complexes Complex – Association of a cation
and an anion or neutral molecule All associated species are dissolved None remain electrostatically effective
Ligand – the anion or neutral molecule that combines with a cation to form a complex Can be various species E.g., H2O, OH-, NH3, Cl-, F-,
NH2CH2CH2NH2
Importance of complexes
Complexing can increase solubility of minerals if ions involved in reactions are complexed Total concentration of species (e.g.,
complexed plus dissolved) will be higher in solution at equilibrium with mineral
E.g., Solution at equilibrium with calcite will have higher SCa2+ if there is also SO4
2- present because of CaSO4o
complex
Some elements more common as complexes Particularly true of metals Cu2+, Hg2+, Pb2+, Fe3+, U4+ usually found
as complexes rather than free ions Their chemical behavior (i.e. mobility,
toxicity, etc) are properties of complex, not the ion
Adsorption affected by complex E.g., Hydroxide complexes of uranyl
(UO22+) readily adsorbed by oxide and
hydroxide minerals OH- and PO4
- complexes readily adsorbed
Carbonate, sulfate, fluoride complexes rarely adsorbed to mineral surfaces
Toxicity and bioavailability depends on complexes Toxicity – e.g. Cu2+, Cd2+, Zn2+, Ni2+,
Hg2+, Pb2+
Toxicity depends on activity and complexes not total concentrations
E.g., CH3Hg+ and Cu2+ are toxic to fish other complexes, e.g., CuCO3
o are not
Bioavailability – some metals are essential nutrients: Fe, Mn, Zn, Cu Their uptake depends on forming
complexes
General observations
Complex stability increases with increasing charge and/or decreasing radius of cation Space issue – length of interactions
Strong complexes form minerals with low solubilities Corollary – Minerals with low solubilities
form strong complexes
High salinity increases complexing More ligands in water to complex
High salinity water increases solubility because of complexing
Complexes – two types
Outer Sphere complexes AKA – “ion Pair”
Inner Sphere complexes AKA – “coordination compounds”
Outer Sphere Complexes Associated hydrated cation and
anion Held by long range electrostatic forces No longer electrostatically effective
Metal ion and ligand still separated by water Association is transient Not strong enough to displace water
surrounding ion Typically smaller ions – Na, K, Ca, Mg, Sr
Larger ions have low charge density Relatively unhydrated Tend to form “contact complexes”
Outer Sphere complexes Metal ion and ligand still separated
by water Association is transient Not strong enough to displace water
surrounding ion Typically smaller ions – Na, K, Ca, Mg, Sr
Larger ions have low charge density Relatively unhydrated Tend to form “contact complexes”
Larger ions have low charge density Relatively unhydrated Tend to form “contact” ion pairs – with
little water in between
Inner Sphere Complexes More stable than ion pairs
Metal and ligands immediately adjacent Metal cations generally smaller than
ligands Largely covalent bonds between metal
ion and electron-donating ligand Charge of metal cations exceeds
coordinating ligands May be one or more coordinating
ligands
Inner sphere – completely oriented water, typically 4 or 6 fold coordination
Outer sphere – partly oriented water
Coordinating cation
An Aquocomplex – H2O is ligand
For ligand, L to form inner-sphere complex Must displace one or more coordinating
waters Bond usually covalent nature E.g.:
M(H2O)n + L = ML(H2O)n-1 + H2O
Size and charge important to number of coordinating ligands: Commonly metal cations smaller than
ligands Commonly metal cation charge exceed
charge on ligands These differences mean cations
typically surrounded by several large coordinating ligands E.g., aquocomplex
Maximum number of ligands depends on coordination number (CN)
Most common CN are 4 and 6, although 2, 3, 5, 6, 8 and 12 are possible
CN depends on radius ratio (RR):
RR = Radius Coordinating Cation
Radius Ligand
Maximum number of coordinating ligands Depends on radius ratio Generates coordination polyhedron
All coordination sites rarely filled Only in aquo-cation complexes
(hydration complexes) Highest number of coordination sites
is typically 3 to 4 The open complexation sites results
from dilute concentration of ligands
Concentrations of solution Water concentrations – 55.6 moles/kg Ligand concentrations 0.001 to 0.0001
mol/kg 5 to 6 orders of magnitude lower
Ligands can bond with metals at one or several sites
Unidentate ligand – contains only one site E.g., NH3, Cl- F- H2O, OH-
Bidentate Two sites to bind: oxalate,
ethylenediamine
Various types
of ligands
Multidentate – several sites for complexing Hexedentate –
ethylenediaminetetraacetic acid (EDTA)
Additional multidentate
ligands
Thermodynamics of complexes
Strength of the complex represented by stability constant Kstab also called Kassociation
An equilibrium constant for formation of complex
Typical metals can form multiple complexes in water with constant composition Al3+, AlF2+, AlF2
+, AlF3
SAl = Al3+ + AlF2+ + AlF2+ + AlF3
Example:Al3+ + 4F- = AlF4
-
Kstab =(aAl3+)(aF-)4
aAlF4
-
Complexation changes “effective concentrations” of solution
Another example:
Ca2+ + SO42- = CaSO4
o
Here the o indicates no charge – a complex
This is not solid anhydrite – only a single molecule
Still dissolved
aCaSO4o is included in the Kstab
calculations It is a dissolved form
Kstab =(aCa2+)(aSO42-)
aCaSO4o
Examples of Kstab calculations and effects of complexing on concentrations