Particle size Ions  molecular clusters  nanocrystals  colloids  bulk minerals Small particles can have a significant % of molecules at their surface Thermodynamics are different (surface free energy) Surface area per mass is huge and charged through interaction with water Sorption of ions to these surfaces can be critical part of contaminant mobility

Surface area Selected mineral groups often occur as colloids / nanoparticles: FeOOH  SA up to 500 m2/g, site density 2-20/nm2 Al(OH)3  SA up to 150 m2/g, site density 2-12/nm2 MnOOH  SA hundreds m2/g, site density 2-20/nm2 SiO2  SA 0.1 – 300 m2/g, site density 4-12/nm2 Clays  SA 10-1000 m2/g, site density 1-5/nm2 Organics  SA up 1300 m2/g, site density 2/nm2

DEFINITIONS Sorption - removal of solutes from solution onto mineral surfaces. Sorbate - the species removed from solution. Sorbent - the solid onto which solution species are sorbed. Three types of sorption: Adsorption - solutes held at the mineral surface as a hydrated species. Absorption - solute incorporated into the mineral structure at the surface. Ion exchange - when an ion becomes sorbed to a surface by changing places with a similarly charged ion previously residing on the sorbent. The three different types of sorption processes defined above cannot always be distinguished clearly in practice. However, it is useful to make these distinctions in theory. When it is not clear exactly which of these processes is occurring, the general term sorption should be used. It should also be kept in mind that not all authors define these processes in exactly the same way as Kehew (2001). Consult Figure 4-27 in Kehew (2001) to make the distinction between adsorption and absorption clearer.

Mineral Surfaces Minerals which are precipitated can also interact with other molecules and ions at the surface Attraction between a particular mineral surface and an ion or molecule due to: Electrostatic interaction (unlike charges attract) Hydrophobic/hydrophilic interactions Specific bonding reactions at the surface

Inner Sphere and Outer Sphere Outer Sphere surface complex  ion remains bounded to the hydration shell so it does not bind directly to the surface, attraction is purely electrostatic Inner Sphere surface complex  ion bonds to a specific site on the surface, this ignores overall electrostatic interaction with bulk surface (i.e. a cation could bind to a mineral below the mineral pHzpc)

Charged Surfaces Mineral surface has exposed ions that have an unsatisfied bond  in water, they bond to H2O, many of which rearrange and shed a H+ ≡S- + H2O  ≡S—H2O  ≡S-OH + H+ OH OH OH2 H+ OH OH OH H+ OH

Surfaces as acid-base reactants The surface ‘SITE’ acts as an amphoteric substance  it can take on an extra H+ or lose the one it has to develop charge ≡S-O- + H+ ↔ ≡S-OH ↔ ≡S-OH2+ The # of sites on a surface that are +, -, or 0 charge is a function of pH pHzpc is the pH where the + sites = - sites = 0 sites and the surface charge is nil OH OH2+ O- OH O- OH OH2+

pHzpc Zero Point of Charge, A.k.a: Zero Point of Net Proton Charge (pHZPNPC) or the Isoelectric Point (IEP) Measured by titration curves (pHzpc similar to pKa…) or electrophoretic mobility (tendency of the solids to migrate towards a positively charged plate) Below pHzpc  more sites are protonated  net + charge Above pHzpc  more sites are unprotonated  net - charge

POINT OF ZERO CHARGE CAUSED BY BINDING OR DISSOCIATION OF PROTONS As can be seen here, the pHpznpc of oxides and silicates varies over a wide range (at least from 2 to > 10). For minerals with very low pHpznpc values (e.g., quartz, feldspars), anion sorption is not likely to be very strong in most natural waters. This is because the pHpznpc value of about 2 is less than the common pH range of natural waters (5.5-8.5), so the surfaces of quartz and feldspars will be negatively charged in most natural waters. On the other hand, minerals with high pHpznpc values (e.g., corundum, Fe-oxides and chrysotile) will generally be more efficient sorbents for anions than for cations, because their surfaces will be positively charged over the pH range of most natural waters.

From Stumm and Morgan, Aquatic Chemistry

ION EXCHANGE REACTIONS Ions adsorbed by outer-sphere complexation and diffuse-ion adsorption are readily exchangeable with similar ions in solution. Cation exchange capacity: The concentration of ions, in meq/100 g soil, that can be displaced from the soil by ions in solution. Also anion exchange capacity for positively charged surfaces

ION EXCHANGE REACTIONS Exchange reactions involving common, major cations are treated as equilibrium processes. The general form of a cation exchange reaction is: nAm+ + mBX  mBn+ + nAX The equilibrium constant for this reaction is given by: In the ion-exchange reaction given above, the X represents the solid surface on which ion exchange occurs.

CATION EXCHANGE CAPACITIES OF MINERALS AND SOILS Note that cation exchange capacities are greatest for 2:1 clay minerals (montmorillonite and vermiculite) and soil organic matter. Oxides, hydroxides and silica sand have the lowest CEC values.

SORPTION ISOTHERMS - I The capacity for a soil or mineral to adsorb a solute from solution can be determined by an experiment called a batch test. In a batch test, a known mass of solid (S m) is mixed and allowed to equilibrate with a known volume of solution (V ) containing a known initial concentration of a solute (C i). The solid and solution are then separated and the concentration (C ) of the solute remaining is measured. The difference C i - C is the concentration of solute adsorbed.

Kd Descriptions of how solutes stick to the surface What would the ‘real’ behavior be you think?? Kd

SORPTION ISOTHERMS - II The mass of solute adsorbed per mass of dry solid is given by where S m is the mass of the solid. The test is repeated at constant temperature but varying values of C i. A relationship between C and S can be graphed. Such a graph is known as an isotherm and is usually non-linear. Two common equations describing isotherms are the Freundlich and Langmuir isotherms.

FREUNDLICH ISOTHERM The Freundlich isotherm is described by where K is the partition coefficient and n  1. When n < 1, the plot is concave with respect to the C axis. When n = 1, the plot is linear. In this case, K is called the distribution coefficient (Kd ). The partition coefficient K is a measure of the degree to which a sorbate partitions between the surface and the solution. The higher the value of K, the greater affinity the sorbate has for the surface. The graph above shows how K and n affect the shape of the Freundlich isotherm. The higher K, the steeper the initial slope of the isotherm. The smaller the value of n, the greater the deviation from linearity (the more concave the isotherm becomes with respect to the C axis. The case where n = 1 does not fit the adsorption of most inorganic solutes. However, a Freundlich isotherm with n = 1 is often used successfully to describe the sorption of hydrophobic organic compounds (e.g., carbon tetrachloride).

LANGMUIR ISOTHERM The Langmuir isotherm describes the situation where the number of sorption sites is limited, so a maximum sorptive capacity (S max) is reached. The governing equation for Langmuir isotherms is: The Langmuir isotherm describes cases in which there are a limited number of sites available for sorption, so the sorption sites become saturated. Note that the Langmuir isotherm has a form very similar to the Michaelis-Menton equation used to describe the kinetics of enzyme-mediated reactions (hyperbolic kinetics). Recall that the Michaelis-Menton equation results from the possibility that saturation of the enzyme may limit the rate of reaction.

Sorption of organic contaminants Organic contaminants in water are often sorbed to the solid organic fractions present in soils and sediments Natural dissolved organics (primarily humic and fulvic acids) are ionic and have a Koc close to zero Solubility is correlated to Koc for most organics

Measuring organic sorption properties Kow, the octanol-water partition coefficient is measured in batches with ½ water and ½ octanol – measures proportion of added organic which partitions to the hydrophobic organic material Empirical relation back to Koc: log Koc = 1.377 + 0.544 log Kow

ADSORPTION OF METAL CATIONS - I In a natural solution, many metal cations compete for the available sorption sites. Experiments show some metals have greater adsorption affinities than others. What factors determine this selectivity? Ionic potential: defined as the charge over the radius (Z/r). Cations with low Z/r release their waters of hydration more easily and can form inner-sphere surface complexes. Cations with low charge to radius ratios (ionic potentials) are not strongly hydrated. These cations can easily shed their waters of hydration to participate in inner-sphere surface complexes. Cations with high ionic potentials are strongly hydrated; they do not surrender their waters of hydration easily, and so are more likely to form outer-sphere surface complexes. Because inner-sphere complexes are stronger than outer-sphere complexes, we would expect that cations with low ionic potentials would sorb more strongly to surfaces than cations with high ionic potentials.

ADSORPTION OF METAL CATIONS - II Many isovalent series cations exhibit decreasing sorption affinity with decreasing ionic radius: Cs+ > Rb+ > K+ > Na+ > Li+ Ba2+ > Sr2+ > Ca2+ > Mg2+ Hg2+ > Cd2+ > Zn2+ For transition metals, electron configuration becomes more important than ionic radius: Cu2+ > Ni2+ > Co2+ > Fe2+ > Mn2+ Experimental results confirm our suspicions: for series of cations with the same oxidation state (isovalent series), larger cations have greater sorption affinities than smaller cations. Because the charge is constant, larger cations have lower ionic potentials than smaller cations. Thus, as expected, lower ionic potentials correlate with higher sorption affinities. For transition metals, there are additional complications. Transition metals, by definition, differ in the number of d-electrons in their valence shells. These different electronic configurations give rise to something called ligand field effects. Ligand field effects are more important than ionic size in determining sorption affinities, resulting in the order given above.

ADSORPTION OF METAL CATIONS - III For variable-charge sorbents, the fraction of cations sorbed increases with increasing pH. For each individual ion, the degree of sorption increases rapidly over a narrow pH range (the adsorption edge). For minerals whose surface charge is determined by variable charge, we find experimentally that the percentage of cations sorbed increases with increasing pH. This is a result of the fact that, at low pH, the mineral surface is positively charged, and tends to repel cations, but at high pH, the surface is negatively charged and tends to attract cations. Each cation exhibits a relatively narrow range of pH (about 2 units) over which its sorption increases from near 0% to near 100%. This is referred to as the adsorption or sorption edge.

Exchange reaction and site competition For a reaction: A + BX = B + AX Plot of log[B]/[A] vs. log[BX]/[AX] yield n and K When n and K=1  Donnan exchange, exhange only dependent on valence, bonding strictly electrostatic When n=1 and K≠1  Simple ion exchange, dependent on valence AND size, bonding strictly electrostatic When n≠1 and K≠1 Power exchange, no physical description (complicated beyond the model) and unbalanced stoichiometry

Electrostatic models Combining electrostatic interactions and specific complexation using mechanistic and atomic ideas about the surface yield models to describe specific sorption behavior