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 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., 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 Outer Sphere complexes Inner 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

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

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

Radius Coordinating Cation RR = Radius Ligand 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 Coordinating Cation RR = 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

Al3+ + 4F- = AlF4- aAlF4- Kstab = (aAl3+)(aF-)4 Typical metals can form multiple complexes in water with constant composition Al3+, AlF2+, AlF2+, AlF3 SAl = Al3+ + AlF2+ + AlF2+ + AlF3 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