Complexes

1. 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

2. 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

3. 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

4. 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

5. 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

6. Bioavailability – some metals are essential nutrients: Fe, Mn, Zn, Cu • Their uptake depends on forming complexes

7. 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

8. High salinity increases complexing • More ligands in water to complex • High salinity water increases solubility because of complexing

9. Complexes – two types • Outer Sphere complexes • AKA – “ion Pair” • Inner Sphere complexes • AKA – “coordination compounds”

10. 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”

11. 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”

12. Larger ions have low charge density • Relatively unhydrated • Tend to form “contact” ion pairs – with little water in between

13. 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

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

15. 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

16. 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

17. 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

18. Maximum number of coordinating ligands • Depends on radius ratio • Generates coordination polyhedron

19. 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

20. 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

21. Ligands can bond with metals at one or several sites • Unidentateligand – contains only one site • E.g., NH3, Cl- F- H2O, OH- • Bidentate • Two sites to bind: oxalate, ethylenediamine

22. Various types of ligands

23. Multidentate – several sites for complexing • Hexedentate – ethylenediaminetetraacetic acid (EDTA)

24. Additional multidentate ligands

25. Thermodynamics of complexes • Strength of the complex represented by stability constant • Kstabalso called Kassociation • An equilibrium constant for formation of complex

26. 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

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

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

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

30. Examples of Kstab calculations and effects of complexing on concentrations