1. IS SOLUBILITY THE ONLY CONTROL ON SOLUTE CONCENTRATIONS?
The answer is NO! Solubility often controls the concentrations of major solutes such as Si, Ca, and Mg, and some minor or trace solutes such as Al and Fe.
However, for many trace elements, sorption processes maintain concentrations below saturation with respect to minerals.
In other words, sorption is a means to remove solutes even when the solution is undersaturated with any relevant solids.
In the preceding two lectures, we have learned about solubility controls on natural water compositions. Both congruent and incongruent dissolution can exert control on the concentrations of major solutes, and some trace solutes as well. Nevertheless, the concentrations of many trace elements are controlled by a collection of processes collectively called sorption. Sorption processes involve the removal of solutes from solution into or onto a solid. These processes may occur even though the solution is not saturated with any mineral containing the solute of interest. For example, the concentration of a trace element such as Cd may be limited by sorption onto the surface of a clay or iron oxyhydroxide mineral, even though the solution is undersaturated with respect to all minerals of which Cd is an essential constituent. Sorption processes are important because they retard the movement of contaminants through aquifers. Sorption processes are expected to play a dominant role in retaining radionuclides near nuclear waste repositories, should the primary waste form be breached and come in contact with ground water. Most repository designs provide for backfilling of metal canisters (containing nuclear waste-bearing borosilicate glass) with clays. Sorption onto clay and other mineral surfaces should help retard the migration of radionuclides into the biosphere. Finally, sorption processes also are important in uncontaminated natural waters (recall the Madison Aquifer example in Lecture 4 where ion exchange occurred along flow path 2). Thus, an understanding of sorption processes is of paramount importance to aqueous geochemistry.

2. 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

3. 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.

4. 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

5. 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+

6. GOUY-CHAPMAN DOUBLE-LAYER MODEL STERN-GRAHAME TRIPLE-LAYER MODEL
This slide compares two different models of a mineral surface in contact with aqueous solutions. The Gouy-Chapman double-layer model is essentially the model we considered in slide 12. The Stern-Grahame triple-layer model applies to situations where the surface potential is so strong (owing to high surface charge) that, in addition to the diffuse ion layer, a compact layer of cations exists immediately adjacent to the mineral surface. The ions in this compact layer are held tightly by electrostatic forces and are not free to move like the ions in the diffuse layer.

7. Sorption to ≡S-OH sites
≡S-OH + M2+  ≡S-OM+ + H+
≡S-OH + L2-  ≡S-L- + OH-
In addition, can also have bi-dendate sorption reactions:
≡S-OH
+ M2+
≡S-O M
+ 2 H+

8. 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

9. 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 ( ), 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.

10. From Stumm and Morgan, Aquatic Chemistry

11. Anion-Cation sorption
Equilibrium description for sorption of:
≡S-OH + M2+  ≡S-OM+ + H+
Where Dz is the stoichiometric net change in surface charge due to the sorption reaction (+1 here), F is Faraday’s constant (96485 Coulombs per mole),  is the electrical potential at the surface, R is the gas constant, and T is temperature in Kelvins, the whole right term is called the coulombic term

12. 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)

13. 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.

14. 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.

15. 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.

16. 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.

17. 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.

18. 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.

19. 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).

20. 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.

21. 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.

22. 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.

23. Organic Geochemistry Organic compounds – where do they come from?
How are they different from inorganic compounds?
What determines if they are reactive (more nonreactive = recalcitrant)

24. 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

25. 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 = log Kow