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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 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 SO42present 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., 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 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) Inner Sphere complexes (stronger bonds) AKA – “ion Pair” 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 An Aquocomplex – H2O is ligand Note – cross section, actually 3-D sphere + Outer sphere – partly oriented water Coordinating cation Inner sphere – completely oriented water, typically 4 or 6 fold coordination For ligand other than water to form innersphere 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, AlF4SAl = Al3+ + AlF2+ + AlF2+ + AlF3 + AlF4- Example: Kstab = Al3+ + 4F- = AlF4aAlF4(aAl3+)(aF-)4 Another example: Ca2+ + SO42- = CaSO4o The o indicates no charge – a complex Since CaSO4º not solid anhydrite –a single molecule Dissolved – must include the CaSO4º in thermodynamic calculations aCaSO4º ≠ mCaSO4º Kstab = aCaSO4o (aCa2+)(aSO42-) Examples of Kstab calculations and effects of complexing on concentrations