Complexes
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Complex – Association of a cation and an
anion or neutral molecule
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All associated species are dissolved
None remain electrostatically effective
Ligand – the anion or neutral molecule
that combines with a cation to form a
complex
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Can be various species
E.g., H2O, OH-, NH3, Cl-, F-, NH2CH2CH2NH2
Importance of complexes
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Complexing can increase solubility of
minerals if ions involved in reactions are
complexed
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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 SO42present because of CaSO4o complex
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Some elements more common as
complexes
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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
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Adsorption affected by complex
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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
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Toxicity and bioavailability depends on
complexes
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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., CuCO3o are not
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Bioavailability – some metals are essential
nutrients: Fe, Mn, Zn, Cu
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Their uptake depends on forming complexes
General observations
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Complex stability increases with increasing
charge and/or decreasing radius of cation
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Space issue – length of interactions
Strong complexes form minerals with low
solubilities
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Corollary – Minerals with low solubilities form
strong complexes
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High salinity increases complexing
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More ligands in water to complex
High salinity water increases solubility
because of complexing
Complexes – two types
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Outer Sphere complexes
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AKA – “ion Pair”
Inner Sphere complexes
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AKA – “coordination compounds”
Outer Sphere Complexes
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Associated hydrated cation and anion
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Metal ion and ligand still separated by
water
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Held by long range electrostatic forces
No longer electrostatically effective
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
Outer Sphere complexes
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Metal ion and ligand still separated by
water
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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
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Relatively unhydrated
Tend to form “contact complexes”
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Larger ions have low charge density
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Relatively unhydrated
Tend to form “contact” ion pairs – with little
water in between
Inner Sphere Complexes
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More stable than ion pairs
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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
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May be one or more coordinating ligands
An Aquocomplex – H2O is ligand
Outer sphere – partly
oriented water
Coordinating cation
Inner sphere – completely
oriented water, typically 4
or 6 fold coordination
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For ligand, L to form inner-sphere complex
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Must displace one or more coordinating
waters
Bond usually covalent nature
E.g.:
M(H2O)n + L = ML(H2O)n-1 + H2O
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Size and charge important to number of
coordinating ligands:
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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
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E.g., aquocomplex
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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
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Maximum number of coordinating ligands
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Depends on radius ratio
Generates coordination polyhedron
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All coordination sites rarely filled
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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
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Concentrations of solution
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Water concentrations – 55.6 moles/kg
Ligand concentrations 0.001 to 0.0001 mol/kg
5 to 6 orders of magnitude lower
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Ligands can bond with metals at one or
several sites
Unidentate ligand – contains only one site
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E.g., NH3, Cl- F- H2O, OH-
Bidentate
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Two sites to bind: oxalate, ethylenediamine
Various
types of
ligands
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Multidentate – several sites for complexing
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Hexedentate – ethylenediaminetetraacetic
acid (EDTA)
Additional
multidentate
ligands
Thermodynamics of complexes
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Strength of the complex represented by
stability constant
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Kstab also called Kassociation
An equilibrium constant for formation of
complex
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Typical metals can form multiple complexes
in water with constant composition
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Al3+, AlF2+, AlF2+, AlF3
SAl = Al3+ + AlF2+ + AlF2+ + AlF3
Example:
Kstab =
Al3+ + 4F- = AlF4aAlF4(aAl3+)(aF-)4
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Complexation changes “effective
concentrations” of solution
Another example:
Ca2+ + SO42- = CaSO4o
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Here the o indicates no charge – a
complex
This is not solid anhydrite – only a single
molecule
Still dissolved
Kstab =
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aCaSO4o
(aCa2+)(aSO42-)
aCaSO4o is included in the Kstab calculations
It is a dissolved form
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Examples of Kstab calculations and effects
of complexing on concentrations