The Geochemistry of Rocks and Natural Waters Course no. 210301 Aquatic Geochemistry A. Koschinsky

1. The Geochemistry of Rocks and Natural WatersCourse no. 210301Aquatic GeochemistryA. Koschinsky

2. Aquatic Geochemistry - Introduction • Some general definitions • Concentration and activity • Molarity (mmol/l) = concentration (mg/l) / atomic weight • Molality = Molarity / density of solution (or seawater, respectively) • Excercise: 1. Express „5 µmol/l Pb“ in the unit mg/l (atomic weight Pb = 207). • 2. Express „60 mg/l CO2“ in the unit mmol/l (molecular weight CO2 = 44). Activity a = c(concentration) x f (activity coefficient, also called ) The activity can be regarded as effective concentration (i.e. the share of ions in solution that is able to react). This form of the extended Debye-Huckel equation is valid for 25°C and solution ionic strength up to 0.1 molar

3. Aquatic Geochemistry - Introduction • Activity coefficients are a function of the following: a. hydrated radius of ion (a) • b. charge on ion (z) • c. ionic strength of • solution (m) • = 1/2 cizi2

4. Aquatic Geochemistry - Introduction • Speciation • Chemical Speciation describes the chemical form in which a compound is present in the system; dissolved species include hydrated free cations and different complex forms. • Importance of Speciation • A. Toxicity and bioavailability of Metalsis related to the form of an ion in solution (e.g., Cu2+ is toxic while CuCO30 is not). • B. Productivity of waters is related to the organic ligand concentration (e.g., organic conditioning in upwelling waters). • C. Solubility of metals is related to the form in solution (e.g., ligands can solubilize metals from sediments).

5. Aquatic Geochemistry - Introduction • Speciation • Factors Controlling the State of Metal Ions • 1. Oxidation State • 2. pH • 3. Composition of Inorganic Ligands (OH-, CO32-, HS-) • 4. Composition of Organic Ligands (Humic and Fulvic Acids, others) • 5. Pressure and Temperature • Importance of pH in Speciation • HX    H+  +  X- • X- is a ligand that can complex a metal (CO32-, OH-, HPO42-, PO43-, organic acid anions)

6. Aquatic Geochemistry - Introduction • Most reactions in aqueous solutions can be placed in one of these categories: • Acid-base reactions, e.g., dissociation of carbonic acid: • H2CO3 H+ + HCO3 • Give the dissociation reaction for another acid: • Complexation, e.g., hydrolysis of mercury: • Hg2+ + H2O  Hg(OH)+ + H+ • Give the complexation reaction for dissolved Zn and chloride: • Dissolution/Precipitation, e.g., dissolution of orthoclase: • KAlSi3O8 + H+ + 7H2O  Al(OH)3 + K+ + 3H4SiO4 • Give the precipitation reaction of calcite: • Adsorption/Desorption, e.g., adsorption of Mn on a clay surface: • =S + Mn2+=S–Mn (where =S indicates the surface of the clay). • Give the adsorption reaction of copper on a goethite surface:

7. Aquatic Geochemistry - Complexation • Ions in solution often associate with other ions, forming complexes. This affects the solubility and reactivity of ions. For example, if ions form stable, soluble complexes, this greatly enhances the solubility of the ions. • Central ion (metal) - ligands (ions or molecules that surround, or coordinate, the central ion) • The simplest complexes are those formed between metals and water: aquo complexes (Water molecules are the ligands). Aquo-complexes are ubiquitous: all charged species will have a solvation shell. Truly “free ions” do not exist. • Example: zinc aquo complex Zn(H2O)62+ Other complexes include: • Inorganic complexes (such as chloride complexes of the form MnCl+, carbonate complexes such as Nd(CO3)+ or hydroxide complexes such as Zr(OH)5+ and Ti(OH)40 • Organic complexes of mostly complicated structure

8. Complexation - Types of complexes • Beyond aquo complexes, we can distinguish two types of complexes: • Ion pairs (outer sphere complexes), where ions of opposite charge associate with one and other through electrostatic attraction, yet each ion retains part or all of it’s solvation sphere. • Complexes (senso stricto, inner sphere complexes), where the two ions are in contact and a bond forms between them that is at least partly covalent in nature.

9. Complexation - Stability Constants Complex formation reaction between a metal cation M and an anion or ligand L : mM + nL <--> MmLn As with any other reaction, the equilibrium constant is defined as: K = [MmLn] / [M]m [L]n Example: Zn(H2O)62+ + OH- <---> Zn(H2O)5(OH)+ + H2O Or: Zn2+ + OH- <---> Zn(OH)+ K1 = [ZnOH+] / [Zn2+] [OH] The zinc ion might associate with a second hydroxyl: Zn(OH)+ + OH- <---> Zn(OH)20 Define the second equilibrium constant K2 = The equilibrium constant for complex formation is often referred to as stability constant.

10. The repulsion between the central ion and the protons depends on: • pH • size of the central ion • charge of the central ion. Complexation - Water-related complexes Aquo-complexes can act as weak acids because the positive charge of the central ion tends to repel hydrogen atoms in the water molecules:

11. Complexation - Other complexes

12. Complexation - Other complexes With respect to complex formation, the elements can be divided into four classes: 1. non-metals (forming anions), ligand formers. 2.“A-type” or “hard” metals. These metals have spherically symmetric, inert-gas type outer electron configurations. Their electron shells are not readily deformed by electric fields and can be viewed as “hard spheres”. They preferentially form complexes with F and ligands having O as the donor atoms (OH–, CO32- , PO43- , SO42-). Stability of the complexes formed by these metals increases with charge to radius ratio. Thus the alkalis form only weak, unstable complexes, while elements such as Zr4+ form very strong, stable complexes. Anions and cations are bound primarily by electrostatic forces, i.e., ionic-type bonds.

13. Complexation - Other complexes B-type, or “soft”, metal ions: Their electron shells are not spherically symmetric and are readily deformed by the electrical fields of other ions (“soft”). They preferentially form complexes with bases having S, I, Br, Cl, or N (such as ammonia) as the donor atom. Bonding between the metal and ligand(s) is primarily covalent and is comparatively strong. Thus Pb form strong complexes with Cl– and S2–. 4. First series transition metals: Their electron sheaths are not spherically symmetric, but they are not so readily polarizable as the B-type metals. On the whole, however, their complex-forming behavior is similar to that of the B-type metals.

14. Complexation - Transition metal complexes Among the transition metals, the sequence of complex stability is Mn2+ < Fe2+ < Co2+ < Ni2+< Cu2+ > Zn2+, a sequence known as the Irving-Williams Series. In the figure, all the sulfate complexes have approximately the same stability, a reflection of the predominately electrostatic bonding between sulfate and metal. Pronounced differences are observed for organic ligands. Although the absolute value of stability complexes varies from ligand to ligand, the relative affinity of ligands having the same donor atom for these metals is always similar.

15. Complexation - Chelation Chelation: Organic molecules can often have more than one functional group and hence can coordinate a metal at several positions, a process called chelation. Such ligands are called multi-dentate and organic compounds having these properties are referred to as chelators or chelating agents (for example EDTA).

16. Complexation - Speciation Calculation

17. Complexation - Speciation Calculation Dependence on pH and temperature

18. 0 0 4 Pb Zn 1 2 4 2 3 3 1 Pb 1 2 Zn 2 3 1 4 0 Complexation - Speciation Calculation Dependance on ligand concentration and temperature

19. Solid - Solution Interactions Mineral surface - interface - solution The processes at the interface govern equilibria between solids and solutions.

20. Solid - solution interactions: Adsorption • Important: Many minerals have large active surface areas (strong sorptive properties)! • Example: In natural marine solids, the following specific surface areas were measured: • Clay fraction of a sediment: 92 m2/g • Manganese nodule (todorokite MnO2) 181 m2/g • Ferromanganese crust (vernadite -MnO2) 350 m2/g • Ferrihydrite FeOOH 140 m2/g • Concentrations of many trace metals in many natural aqueous systems are not controlled by precipitation/dissolution, but by adsorption/desorption on/from mineral and organic surfaces. • Solution complexation and adsorption are competing reactions in aquatic systems!

21. Solid - solution interactions: Adsorption Equilibria between dissolved and surface-bound species • Equilibration in dissolved milieu: formation of dissolved complexes by central ion and ligands • Stability constant  of the complex • - Distribution of compounds between seawater and solid material (particulate suspendended matter, bottom sediment, organisms) • Distribution or partitioning coefficient KdMe(diss.)<----> Me(solid) - Competition between reaction of the ion with a ligand (solution complexation) and reaction of the ion with a solid surface (sorption)

22. Adsorption mechanisms • Adsorption: attachment of an ion in solution to a preexisting solid surface (clay, oxide, carbonate, … particle). Adsorption involves: • Electrostatic interactions: Solid surfaces are typically electrically charged. This electrostatic force, which is effective over greater distances than purely chemical forces, affects surface complex formation and loosely binds other ions to the surface.

23. Adsorption mechanisms • Surface complex formation: The formation of coordinative bonds between metals and ligands at the surface (similar to the formation of complexes between dissolved components). • Hydrophobic adsorption: Many organic substances are highly insoluble in water due to their non-polar nature. These substances become adsorbed to surfaces because they are repelled by water.

24. Adsorption mechanisms As is the case with soluble complexes, surface complexes may be divided into inner sphere and outer sphere complexes. Inner sphere complexes involve some degree of covalent bonding between the adsorbed species and atoms on the surface. In an outer-sphere complex, one or more water molecules separate the adsorbed ion and the surface; in this case adsorption involves only electrostatic forces. The third possibility is that an ion may be held within the diffuse layer by long-range electrostatic forces.

25. - - - - - -H - - - -H - -H Oxide + H+ --> Oxide + H+ --> Oxide … Adsorption mechanisms The functional groups of mineral and organic surfaces have properties similar to those of their dissolved counterparts. In this sense, surface complexation reactions are similar to complexation reactions in solution. Functional groups: -OH, =O (inorganic), -COOH (carboxyl group, organic acid), … Reactions between these surface groups and dissolved species are influenced by the proximity of surface groups to each other. For example, the surface charge will change systematically as the adsorbed surface concentration of a positive species such as H+ increases. This change in surface charge will decrease the attraction between H+ ions and the surface. As a result, the equilibrium constant for the surface protonation reaction will change as the surface concentration of H+ increases. Equilibrium constant or distribution coefficient Kd = [Csorbed]/[Caq]

26. Outer- and inner-sphere complexes (pK = -logK)

27. Surface complexation Stable surface complexes are most likely to be formed at surface irregularities such as kinks and steps (see scheme below, formation of copper complexes on calcite surface). The presence of other metal ions may lead to the incorporation of these ions into the crystal lattice, or to the inhibition of the calcite crystal growth. Both processes involve the substitution of Ca by these ions at the surface complexation sites.

28. Oxygen and metal atoms at an oxide surface are incompletely coordinated; i.e., they are not surrounded by oppositely charged ions as they would be in the interior of a crystal (Fig. a). Consequently, mineral surfaces immersed in water bind water molecules (Fig. b), which can then dissociate, leaving a hydroxyl group bound to the surface metal ion: In a similar fashion, incompletely coordinated oxygens at the surface can also bind water molecules, which can then dissociate, again creating a surface hydroxyl group: Thus the surface on an oxide immersed in water quickly becomes covered with hydroxyl groups (Fig. c), which are considered to constitute part of the surface rather than the solution. Adsorption: oxide surfaces in water

29. These hydroxyl groups can then act as either proton acceptors or proton donors through further association or dissociation reactions, e.g.: These reactions are strongly pH-dependent! Adsorption: oxide surfaces in water

30. 2- Adsorption: oxide surfaces in water Adsorption on a mineral surface may occur when a metal replaces a surfaceproton, or (b) a ligand replaces a surface OH group. The adsorbed metal (c) may bind an additional ligand, and the ligand (d) may bind an additional metal. Multidentate adsorption involves more than one surface site (e, f).

31. pH dependence of adsorption Since surface bound protons and OH groups are almost inevitably involved in adsorption, adsorption of metals and ligands is strongly pH dependent. This strong dependence on pH certainly reflects protonation of the surface, but it also reflects the extent of hydrolysis of the ion in solution.

32. Surface charge of particles • Mineral surfaces develop electrical charge for several reasons: • Complexation reactions between the surface and dissolved species, including protonation and deprotonation. Because these reactions depend on pH, this aspect of surface charge is pH dependent. • Lattice imperfections at the solid surface as well as substitutions within the crystal lattice (e.g., Al3+ for Si4+). Because the ions in interlayer sites of clays are readily exchangeable, this mechanism is particularly important in the development of surface charge in clays. • Thus there are several contributions to surface charge density. We define net as the net density of electric charge on the solid surface, and can express it as: • net = 0 + H + SC • where 0 is the intrinsic surface charge due to lattice imperfections and substitutions, • H is the net proton charge, i.e., the charge due to binding H+ and OH–, • SC is the charge due to other surface complexes. •  is usually measured in coulombs per square meter (C/m2).

33. Surface charge of particles Surface charge of some common sedimentary materials as a function of pH. The pH dependence of surface charge reflects the predominance of attached protons at low pH and the predominance of attached hydroxyls at higher pH.

34. 2 2 Surface charge of particles At some value of pH the surface charge will be zero. This point is known as the isoelectric point, or zero point of charge (ZPC). This occurs when the charge due to adsorption of cations is balanced by charge due to adsorption of anions. When the pHzpc of a solid is lower than the surrounding pH, then the solid develops a negative surface charge. When the pHzpc of a solid is higher than the surrounding pH, then the solid develops a positive surface charge. Question: What is the surface charge of goethite and vernadite (-MnO2) in seawater?

35. Long-Term Fate of Adsorbed Ions • Desorption of ions is much less favored than sorption. • Sorbed ion complexes may polymerize to form surface precipitatites. • Sorbed ions may be incorporatedinto the host structure via dissolution and recrystallization.

36. Surface precipitation and solid solution In surface precipitation, cations (or anions) which adsorb to the surface of a mineral, may form a precipitate of the cation (anion) with the constituent ions of the mineralat high surface coverage. Cations at the solid-water interface are treated as surface species, while those not in contact with the solution phase are treated as solid species forming a solid solution. The formation of solid solution implies isomorphic substitution. Surface precipitate Solid solution

37. Surface precipitation and solid solution There is a continuum of surface complexation (adsorption) and surface precipitation. At low sorbate concentrations, surface complexation is the dominant mechanism (Fig. a). As the sorbate concentration increases, surface precipitation becomes the dominant “sorption” (metal ion incorporation) mechanism (Fig. b).

38. Aquatic Geochemistry - Precipitation/Dissolution • A few general rules: • Precipitation of ions as insoluble (low solubility) minerals is a major mechanism by which metal concentrations are limited in groundwater and other natural waters. • A useful generalization is that highly charged cations (+2, +3, +4) form insoluble precipitates with highly charged anions (sulfate, carbonate, phosphate) • Chalcophile elements (Pb, Cd, Cu, Zn, Hg) form insoluble sulfides.

39. Dissolution/precipitation - Some definitions

40. Dissolution/precipitation - carbonate minerals

41. Dissolution/precipitation - hydroxide minerals

42. Dissolution/precipitation - sulfate minerals

43. Dissolution/precipitation - sulfide minerals

44. Dissolution/precipitation - solubility calculations 8 Convert the result in mol/l to ppm!

45. Dissolution/precipitation - solubility calculations than PbSO4 The presence of minerals such as CaCO3 and CaSO4 can suppress the solubility of metals such as Pb and Cd via the common ion effect.

46. Dissolution/precipitation - solubility calculations 2,7,9 ? Because they are strong bases, the solubilities of hydroxides, sulfides and carbonates will depend on the pH.