the geochemistry of natural waters

1: 1 THE GEOCHEMISTRY OF NATURAL WATERS MINERAL WEATHERING AND MINERAL SURFACE PROCESSES - I CHAPTER 4 - Kehew (2001) Silica solubility

2: 2 We will spend three lectures on Chapter 4 in Kehew (2001). In this lecture, we will cover four main topics. First, we will discuss some factors that govern the rates of weathering of silicate and oxide minerals, including the relative stabilities of primary minerals as given by Goldich’s series, and the nature of incongruent dissolution. Next we will briefly discuss the nature of the most common solid products of weathering, i.e., clay minerals. We will then briefly discuss the mechanisms by which oxides and silicates dissolve. Finally, we will investigate somewhat in depth the thermodynamics of dissolution of silica, and Al and Fe oxyhydroxides. We will spend three lectures on Chapter 4 in Kehew (2001). In this lecture, we will cover four main topics. First, we will discuss some factors that govern the rates of weathering of silicate and oxide minerals, including the relative stabilities of primary minerals as given by Goldich’s series, and the nature of incongruent dissolution. Next we will briefly discuss the nature of the most common solid products of weathering, i.e., clay minerals. We will then briefly discuss the mechanisms by which oxides and silicates dissolve. Finally, we will investigate somewhat in depth the thermodynamics of dissolution of silica, and Al and Fe oxyhydroxides.

3: 3 SILICA SOLUBILITY - I In the absence of organic ligands or fluoride, quartz solubility is relatively low in natural waters. Below pH 9, the dissolution reaction is: SiO2(quartz) + 2H2O(l) ? H4SiO40 for which the equilibrium constant at 25°C is: At pH < 9, quartz solubility is independent of pH. Quartz is frequently supersaturated in natural waters because quartz precipitation kinetics are slow. With respect to solubility, everything is relative. The solubility of silica is low relative to a salt like NaCl, but is high compared to aluminum and iron oxyhydroxides. If ligands that can form strong complexes with silicon, e.g., fluoride or organic ligands (the role of organic ligands in silica solubility is somewhat controversial) are absent, then the main form of dissolved silica in most natural waters is silicic acid or H4SiO40. As indicated by the reaction given in this slide, the concentration of silicic acid in equilibrium with quartz (one form of silica) is independent of pH. Because H4SiO40 is the predominant form of silica in solutions with pH < 9, the solubility of quartz is therefore independent of pH under these conditions. The solubility of silica is also not strongly dependent on the ionic strength, because activity coefficients of neutral species are very close to unity (see Lecture 2). However, because the equilibrium constant for the reaction shown above is dependent on temperature and pressure, silica solubility is also dependent on these factors. In fact, in geothermal systems, the solubility of silica can be used as a geothermometer to determine the temperature at which the fluid last reached equilibrium with quartz or amorphous silica. Quartz precipitation kinetics at low temperature are quite low. On the other hand, concentrations of H4SiO40 may build up to quite high values during the weathering of aluminosilicate minerals. It is often found that H4SiO40 concentrations in natural waters exceed those dictated by quartz solubility. In other words, many natural waters are supersaturated with quartz. With respect to solubility, everything is relative. The solubility of silica is low relative to a salt like NaCl, but is high compared to aluminum and iron oxyhydroxides. If ligands that can form strong complexes with silicon, e.g., fluoride or organic ligands (the role of organic ligands in silica solubility is somewhat controversial) are absent, then the main form of dissolved silica in most natural waters is silicic acid or H4SiO40. As indicated by the reaction given in this slide, the concentration of silicic acid in equilibrium with quartz (one form of silica) is independent of pH. Because H4SiO40 is the predominant form of silica in solutions with pH < 9, the solubility of quartz is therefore independent of pH under these conditions. The solubility of silica is also not strongly dependent on the ionic strength, because activity coefficients of neutral species are very close to unity (see Lecture 2). However, because the equilibrium constant for the reaction shown above is dependent on temperature and pressure, silica solubility is also dependent on these factors. In fact, in geothermal systems, the solubility of silica can be used as a geothermometer to determine the temperature at which the fluid last reached equilibrium with quartz or amorphous silica. Quartz precipitation kinetics at low temperature are quite low. On the other hand, concentrations of H4SiO40 may build up to quite high values during the weathering of aluminosilicate minerals. It is often found that H4SiO40 concentrations in natural waters exceed those dictated by quartz solubility. In other words, many natural waters are supersaturated with quartz.

4: 4 SILICA SOLUBILITY - II Thus, quartz saturation does not usually control the concentration of silica in low-temperature natural waters. Amorphous silica can control dissolved Si: SiO2(am) + 2H2O(l) ? H4SiO40 for which the equilibrium constant at 25°C is: Quartz is formed diagenetically through the following sequence of reactions: opal-A (siliceous biogenic ooze) ? opal-A’ (nonbiogenic amorphous silica) ? opal-CT ? chalcedony ? microcrystalline quartz Quartz saturation generally does not control silica concentrations in natural waters. The precipitation rate of amorphous silica is faster than that of quartz, so natural waters are rarely supersaturated with amorphous silica. Natural waters may be saturated with amorphous silica. In other words, the solubility of amorphous silica can control the concentration of dissolved silica in natural waters. According to the equilibrium constants given in this slide and slide 3, the solubility of amorphous silica is approximately 20 times greater than that of quartz. Quartz can precipitate directly from natural waters at low temperatures, but because the rate of this reaction is slow, it is more common for amorphous silica to precipitate first. Because quartz has a lower solubility than amorphous silica, it is the stable phase and, with time, the amorphous silica slowly transforms to quartz. It has been shown that, during the diagenesis of marine sediments, the sequence shown in this slide is followed. Quartz saturation generally does not control silica concentrations in natural waters. The precipitation rate of amorphous silica is faster than that of quartz, so natural waters are rarely supersaturated with amorphous silica. Natural waters may be saturated with amorphous silica. In other words, the solubility of amorphous silica can control the concentration of dissolved silica in natural waters. According to the equilibrium constants given in this slide and slide 3, the solubility of amorphous silica is approximately 20 times greater than that of quartz. Quartz can precipitate directly from natural waters at low temperatures, but because the rate of this reaction is slow, it is more common for amorphous silica to precipitate first. Because quartz has a lower solubility than amorphous silica, it is the stable phase and, with time, the amorphous silica slowly transforms to quartz. It has been shown that, during the diagenesis of marine sediments, the sequence shown in this slide is followed.

5: 5 SILICA SOLUBILITY - III At pH > 9, H4SiO40 dissociates according to: H4SiO40 ? H3SiO4- + H+ H3SiO4- ? H2SiO42- + H+ The total solubility of quartz (or amorphous silica) is: Being an acid, H4SiO40 can dissociate at elevated pH. The value of pK1 = 9.9 suggests that it is a very weak acid, and that it will only undergo significant dissociation at pH > 9. At pH = 9.9, H4SiO40 and H3SiO4- are present in equal amounts, but at pH > 9.9, the latter predominates. The pK2 value of 11.7 indicates that at pH > 11.7, H2SiO42- becomes the predominant species. The total solubility of silica is the sum of all silica species in solution. Because the concentrations of H3SiO4- and H2SiO42- pH-dependent, once these species become predominant over H4SiO40, silica solubility also becomes pH-dependent. It should be kept in mind that, as long as quartz or amorphous silica is present, and the solution remains in equilibrium with one of these phases, then the concentration of H4SiO40 remains constant, even though this species tends to dissociate to H3SiO4- and H2SiO42- as the pH rises. As some of the H4SiO40 dissociates, more quartz or amorphous silica dissolves to replace the H4SiO40 lost to dissociation. Because H4SiO40 is constant, significant dissociation leads to increased total silica in solution, because eventually H3SiO4- and H2SiO42- make important contributions to dissolved silica on top of the constant amount of H4SiO40 always present. Being an acid, H4SiO40 can dissociate at elevated pH. The value of pK1 = 9.9 suggests that it is a very weak acid, and that it will only undergo significant dissociation at pH > 9. At pH = 9.9, H4SiO40 and H3SiO4- are present in equal amounts, but at pH > 9.9, the latter predominates. The pK2 value of 11.7 indicates that at pH > 11.7, H2SiO42- becomes the predominant species. The total solubility of silica is the sum of all silica species in solution. Because the concentrations of H3SiO4- and H2SiO42- pH-dependent, once these species become predominant over H4SiO40, silica solubility also becomes pH-dependent. It should be kept in mind that, as long as quartz or amorphous silica is present, and the solution remains in equilibrium with one of these phases, then the concentration of H4SiO40 remains constant, even though this species tends to dissociate to H3SiO4- and H2SiO42- as the pH rises. As some of the H4SiO40 dissociates, more quartz or amorphous silica dissolves to replace the H4SiO40 lost to dissociation. Because H4SiO40 is constant, significant dissociation leads to increased total silica in solution, because eventually H3SiO4- and H2SiO42- make important contributions to dissolved silica on top of the constant amount of H4SiO40 always present.

6: 6 SILICA SOLUBILITY - IV The equations for the dissociation constants of silicic acid can be rearranged (assuming a = M ) to get: We can now write: To calculate the concentrations of H3SiO4- and H2SiO42- we need to rearrange the mass-action expressions for the dissociation reactions of silicic acid as shown in this slide. For simplicity we assume that activity coefficients are equal to unity. The expressions we derive for the concentrations of these species turn out to be dependent on the concentration of H4SiO40, but we have already demonstrated that this is a constant at fixed temperature and pressure, if the solution is in equilibrium with either quartz or amorphous silica. Thus, we see that the concentrations of H3SiO4- and H2SiO42- in equilibrium with quartz or amorphous silica are dependent on the activity of hydrogen ion. If we take the logarithm of both sides of the first two equations in this slide, employ the definition of pH, and rearrange the equations a bit, we obtain log MH3SiO4- = log (K1MH4SiO40) + pH and log MH2SiO42- = log (K1K2MH4SiO40) + 2pH To summarize these results, the concentration of H4SiO40 is independent of pH, the concentration of H3SiO4- increases one log unit for each unit increase in pH, and the concentration of H2SiO42- increases two log units for each unit increase in pH. If we plotted the logarithm of the concentrations of each of these species vs. pH, we would get a horizontal line for H4SiO40, a line with slope +1 for H3SiO4-, and a line with slope +2 for H2SiO42-. The slopes of the lines for the concentrations of these species will be the same irrespective of whether the solution is saturated with quartz or amorphous silica. However, the lines will all be shifted vertically for amorphous silica compared to quartz, because the former is the more soluble phase. To calculate the concentrations of H3SiO4- and H2SiO42- we need to rearrange the mass-action expressions for the dissociation reactions of silicic acid as shown in this slide. For simplicity we assume that activity coefficients are equal to unity. The expressions we derive for the concentrations of these species turn out to be dependent on the concentration of H4SiO40, but we have already demonstrated that this is a constant at fixed temperature and pressure, if the solution is in equilibrium with either quartz or amorphous silica. Thus, we see that the concentrations of H3SiO4- and H2SiO42- in equilibrium with quartz or amorphous silica are dependent on the activity of hydrogen ion. If we take the logarithm of both sides of the first two equations in this slide, employ the definition of pH, and rearrange the equations a bit, we obtain log MH3SiO4- = log (K1MH4SiO40) + pH and log MH2SiO42- = log (K1K2MH4SiO40) + 2pH To summarize these results, the concentration of H4SiO40 is independent of pH, the concentration of H3SiO4- increases one log unit for each unit increase in pH, and the concentration of H2SiO42- increases two log units for each unit increase in pH. If we plotted the logarithm of the concentrations of each of these species vs. pH, we would get a horizontal line for H4SiO40, a line with slope +1 for H3SiO4-, and a line with slope +2 for H2SiO42-. The slopes of the lines for the concentrations of these species will be the same irrespective of whether the solution is saturated with quartz or amorphous silica. However, the lines will all be shifted vertically for amorphous silica compared to quartz, because the former is the more soluble phase.

7: 7 This plot illustrates the principles discussed in the previous slide. The light red lines show the concentrations of the various dissolved silica species. The concentration of H4SiO40 is represented by the horizontal line. The concentration of H3SiO4- is represented by the straight line with slope +1, and that of H2SiO42- by the straight line with slope +2. The points where the lines for the concentrations of two successive species cross occur at the pK values for silicic acid. For example, the lines representing the species H4SiO40 and H3SiO4- cross at pH = pK1 = 9.9, and the lines for the species H3SiO4- and H2SiO42- cross at pH = pK2 = 11.7. The heavy dark red curve represents the logarithm of the sum of the concentrations of all the species, that is, the total solubility of quartz. At pH < 9.9, H4SiO40 accounts for almost all the dissolved silica, so the curve representing the total solubility is nearly coincident with the line representing the concentration of H4SiO40. Similarly, at 9.9 < pH < 11.7, the predominant species is H3SiO4-, so the total solubility curve has a slope near +1, and at pH > 11.7, the total solubility curve follows the line representing the concentration of H2SiO42-. Near each of the pK values (i.e., the crossover points), significant and nearly equal concentrations of two species are present, so the total solubility curve rises above the lines for the individual species. Also shown on this plot is the total solubility curve for amorphous silica (dotted green line). This curve has exactly the same shape as that for quartz, but is displaced upward by 1.3 log units (log 20 = 1.3), which is reflective of the fact that amorphous silica is 20 times more soluble than quartz. The plot illustrates that, over the pH range of most natural waters, silica solubility is independent of pH. However, as pH rises above 9, the solubility of silica can increase dramatically. This plot illustrates the principles discussed in the previous slide. The light red lines show the concentrations of the various dissolved silica species. The concentration of H4SiO40 is represented by the horizontal line. The concentration of H3SiO4- is represented by the straight line with slope +1, and that of H2SiO42- by the straight line with slope +2. The points where the lines for the concentrations of two successive species cross occur at the pK values for silicic acid. For example, the lines representing the species H4SiO40 and H3SiO4- cross at pH = pK1 = 9.9, and the lines for the species H3SiO4- and H2SiO42- cross at pH = pK2 = 11.7. The heavy dark red curve represents the logarithm of the sum of the concentrations of all the species, that is, the total solubility of quartz. At pH < 9.9, H4SiO40 accounts for almost all the dissolved silica, so the curve representing the total solubility is nearly coincident with the line representing the concentration of H4SiO40. Similarly, at 9.9 < pH < 11.7, the predominant species is H3SiO4-, so the total solubility curve has a slope near +1, and at pH > 11.7, the total solubility curve follows the line representing the concentration of H2SiO42-. Near each of the pK values (i.e., the crossover points), significant and nearly equal concentrations of two species are present, so the total solubility curve rises above the lines for the individual species. Also shown on this plot is the total solubility curve for amorphous silica (dotted green line). This curve has exactly the same shape as that for quartz, but is displaced upward by 1.3 log units (log 20 = 1.3), which is reflective of the fact that amorphous silica is 20 times more soluble than quartz. The plot illustrates that, over the pH range of most natural waters, silica solubility is independent of pH. However, as pH rises above 9, the solubility of silica can increase dramatically.

8: 8 SILICA SOLUBILITY - V An alternate way to understand quartz solubility is to start with: SiO2(quartz) + 2H2O(l) ? H4SiO40 Now adding the two reactions: SiO2(quartz) + 2H2O(l) ? H4SiO40 Kqtz H4SiO40 ? H3SiO4- + H+ K1 SiO2(quartz) + 2H2O(l) ? H3SiO4- + H+ K We can derive the relationships required to make the preceding diagram in an alternate fashion. This approach depends on the fact that, if we add two chemical reactions, the equilibrium constant of the resulting reaction is equal to the product of the equilibrium constants of the two reactions that were added together. As before, we see that the concentration of H4SiO40 in water in equilibrium with quartz is a constant, independent of pH. We can derive the relationships required to make the preceding diagram in an alternate fashion. This approach depends on the fact that, if we add two chemical reactions, the equilibrium constant of the resulting reaction is equal to the product of the equilibrium constants of the two reactions that were added together. As before, we see that the concentration of H4SiO40 in water in equilibrium with quartz is a constant, independent of pH.

9: 9 SILICA SOLUBILITY - VI Taking the log of both sides and rearranging we get: Finally adding the three reactions: SiO2(quartz) + 2H2O(l) ? H4SiO40 Kqtz H4SiO40 ? H3SiO4- + H+ K1 H3SiO4- ? H2SiO42- + H+ K2 SiO2(quartz) + 2H2O(l) ? H2SiO42- + 2H+ K The manipulations shown in this slide are pretty self-explanatory. Make sure that you know how to derive the plot in slide 7 by at least one of the two methods shown. The manipulations shown in this slide are pretty self-explanatory. Make sure that you know how to derive the plot in slide 7 by at least one of the two methods shown.

10: 10 SILICA SOLUBILITY - VII SUMMARY Silica solubility is relatively low and independent of pH at pH < 9 where H4SiO40 is the dominant species. Silica solubility increases with increasing pH above 9, where H3SiO4- and H2SiO42- are dominant. Fluoride, and possibly organic compounds, may increase the solubility of silica. Saturation with quartz does not control silica concentrations in low-temperature natural waters; saturation with amorphous silica may.