Complexes Complex – Association of a cation

and an anion or neutral molecule All associated species are dissolved None remain electrostatically effective

Importance of complexes

Complexing can increase solubility of minerals if ions involved in reactions are complexed Total concentration (SE) = complexed

plus dissolved Total concentration is higher in solution

than 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

Importance of Complexes

Some elements more common as complexes Particularly true of metals Cu2+, Hg2+, Pb2+, Fe3+, U4+ usually found

as complexes rather than free ions The chemical behavior depends on

complex, not ion, e.g.: Mobility Bioreactivity: Toxicity & bioavialability

Mobility

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

Bioreactivity

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 High charge = stronger bond

Strong complexes form minerals with low solubilities Corollary – Minerals with high solubilities

form weak complexes

High salinity increases complexing More ligands in water to complex

High salinity water increases solubility because of complexing

Complexes – two types

No consistent nomenclature Outer Sphere complexes (weaker

bonds) AKA – “ion Pair”

Inner Sphere complexes (stronger bonds) AKA – “coordination compounds” AKA – “complex” (S&M)

These are ideal end-members – most complexes intermediate in structure

Outer Sphere Complexes Associated hydrated (usually) cation

and anion Held by long range electrostatic forces Fairly weak complex, but ions still no

longer “electrostatically effective” Separated by water molecules

oriented around cation Water separates ions making up

complex

Outer Sphere complexes Association is transient Not strong enough to displace water

surrounding ion Typically smaller cations

Na, K - monovalent so weaker bonds Ca, Mg, Sr - divalent so stronger bonds

Outer Sphere complexes Also larger ions (Cs & Rb) have low

charge density Relatively unhydrated Tend to form “contact complexes” – e.g.,

no water separation Still considered ion pairs, but no

intervening water

Inner Sphere Complexes More stable than ion pairs Form with ligands 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

Inner Sphere Complexes 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

+

Note – cross section, actually 3-D sphere

For ligand other than water 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 A good example is the “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, AlF4-

SAl = Al3+ + AlF2+ + AlF2+ + AlF3 + AlF4

-

Example:Al3+ + 4F- = AlF4

-

Kstab =(aAl3+)(aF-)4

aAlF4-

Another example:

The o indicates no charge – a complex

Ca2+ + SO42- = CaSO4

o

Since CaSO4º not solid anhydrite –a single molecule Dissolved – must include the CaSO4º in

thermodynamic calculations aCaSO4º ≠ mCaSO4º

Kstab =(aCa2+)(aSO42-)

aCaSO4o

Examples of Kstab calculations and effects of complexing on concentrations