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 SO42- 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 In general: Carbonate, sulfate, fluoride complexes rarely adsorbed to mineral surfaces OH- and PO4- complexes readily adsorbed

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., CuCO3o 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 Inner Sphere complexes Outer Sphere complexes AKA – “coordination compounds” Outer Sphere complexes AKA – “ion Pair”

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 Typically smaller ions – Na, K, Ca, Mg, Sr Larger ions have low charge density Relatively unhydrated Tend to form “contact complexes”

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

An Aquocomplex – H2O is ligand Outer sphere – partly oriented water Coordinating cation Inner sphere – completely oriented water, typically 4 or 6 fold coordination

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

Radius Cation RR = Radius Anion 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): Radius Cation RR = Radius Anion

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-

Multidentate – several sites for complexing Bidentate: oxalate, ethylenediamine Hexedentate – ethylenediaminetetraacetic acid (EDTA)

Strength of the compound represented by stability constant Kstab also called Kassociation An equilibrium constant for formation of complex

Al3+ + 4F- = AlF4- aAlF4- Kstab = (aAl3+)(aF-)4 Typical metals form multiple complexes in a single water Al3+, AlF2+, AlF2+, AlFe3 SAl = Al3+ + AlF2+ + AlF2+ + AlFe3 Example: Al3+ + 4F- = AlF4- aAlF4- Kstab = (aAl3+)(aF-)4

Complexation changes “effective concentrations” of solution Another example: Ca2+ + SO42- = CaSO4o

Here the o indicates no charge – a complex This is not solid anhydrite – only a single molecule Still dissolved

Kstab = (aCa2+)(aSO42-) aCaSO4o is included in the Kstab calculations It is a dissolved form

Examples of Kstab calculations and effects of complexing on concentrations