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Dissociation and pH

Dissociation and pH. Dissociation of weak acids/bases controlled by pH Knowing the total amount of S and pH , we can calculate activities of all species and generate curves Example: H 2 S. Hydrogen Sulfide Activity Diagram. Hydrogen Sulfide Activity Diagram. Solubility of Quartz.

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Dissociation and pH

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  1. Dissociation and pH • Dissociation of weak acids/bases controlled by pH • Knowing the total amount of S and pH, we can calculate activities of all species and generate curves • Example: H2S

  2. Hydrogen Sulfide Activity Diagram

  3. Hydrogen Sulfide Activity Diagram

  4. Solubility of Quartz • The oxides of many metals react with H2O to form bases • SiO2(s)+ 2H2O H4SiO4°

  5. Quartz Activity Diagram • When including a solid, the activity diagram looks a little different • Showing fields of stability for each species • Note: we don’t need to define initial log[SiO2] concentration • Activity of solid = 1

  6. Quartz Activity Diagram

  7. H4SiO4

  8. Buffering of pH • Weak acids and bases can buffer pH of a solution • pH changes very little as acid (or base) is added • Need both a protonated and unprotonated species present in significant concentrations • e.g., H2CO3(aq) and HCO3- • Carbonic acid-bicarbonate is the major buffer in most natural waters • Organic acids and sometimes silicic acid can be important buffers

  9. pH Buffering capacity of an aquifer: Minerals as well as aqueous species • Reactions with minerals: carbonate most important, fastest • CaCO3 + H+ ↔ Ca2+ + HCO3- • Silicates, slower, less important • 2KAlSi3O8 + 2H2CO3+ 9H2O  Al2Si2O5(OH)4 + 2K++ 4H4SiO4 + 2HCO3- • H2CO3 consumes acid, HCO3- creates alkalinity • Ion exchange of charge surfaces • Negatively charged S- + H+ ↔ SH

  10. Dissolved Inorganic Carbon (DIC) • Initially, DIC in groundwater comes from CO2 • CO2(g) + H2O ↔ H2CO3° • Equilibrium expression with a gas is known as Henry’s Law • PCO2: partial pressure (in atm or bar); pressure in atmosphere exerted by CO2 • Assuming atmospheric pressure of 1 atm, PCO2 = 10-3.5; concentration of CO2 = 350 ppm • At atm = 1, N2 is 78%, PN2 = 0.78, O2 21%, PO2 = 0.21

  11. Dissolved Inorganic Carbon (DIC) • PCO2of soil gas can be 10-100 times the PCO2 of atmosphere • PCO2 for surface water controlled by atmosphere and biological processes • Photosynthesis (day): drives PCO2 down, less H2CO3, pH increases • 6CO2 + 6H2O + Energy ↔ C6H12O6 + 6O2 • Respiration: increases PCO2, more H2CO3, pH drops

  12. Dissolved Inorganic Carbon (DIC) • In groundwater, no photosynthesis, no diurnal variations • CO2 usually increases along a flow path due to biodegradation in a closed system • CH2O + O2 CO2 + H2O • CH2O = generic organic matter

  13. DIC and pH in Open System • CO2 can be dissolved into or volatilize out of water freely • Surface waters • PCO2 is constant = 10-3.5atm at Earth’s surface

  14. DIC and pH in Open System • What is the pH of natural rainwater? • Controlled by DIC equilibrium • At 25°C, KCO2 = 10-1.47

  15. DIC and pH in Closed System • In a closed system (no CO2 exchange), for a given amount of TIC, speciation is a function of pH • CO2 + H2O ↔ H2CO3 ↔ HCO3- + H+ ↔ CO32- + H+ • At pH = 6.35, [H2CO3] = [HCO3-] • At pH = 10.33, [HCO3-] = [CO32-] • We can do same calculations we did for H2S

  16. Total DIC = 10-1 M pH = 10.33 pH = 6.35 Common pH range in natural waters

  17. Rainwater pH and PCO2 • What if we double PCO2 (10-1.75atm) • [H2CO3] = [10-1.47] [10-1.75] = 10-3.22 • Doubling the PCO2 does not have a large effect on pH • Acid rain can have pH < 4 • Due to other acids (nitric and sulfuric) that are injected into the atmosphere by vehicles and smokestacks

  18. Special points about DIC, pH, and other weak acids • At pH 6.35, Ka1 = [H+], therefore [H2CO3] = [HCO3-] • Likewise, at pH 10.33, Ka2 = [H+], therefore [HCO3-] = [CO32-]

  19. Special points about DIC, pH, and other weak acids • When pH = pKa, concentration of protonated in reactant = deprotonated in product • pKa = -log Ka • for H2CO3 ↔ HCO3- + H+, Ka = 10-6.35, pKa = 6.35 • so for H4SiO4 ↔ H3SiO4- + H+, pKa= 9.71 • And for H3SiO4- H+ + H2SiO42-, pKa= 13.28

  20. Alkalinity • Alkalinity = acid neutralizing capability (ANC) of water • Total effect of all bases in solution • Typically assumed to be directly correlated to HCO3-concentration in groundwater

  21. Alkalinity • Total alkalinity = [HCO3-] + 2[CO32-] + [B(OH) 4-] + [H3SiO4-] + [HS-] + [OH-] – [H+] • Typically in groundwater, [HCO3-] >> [CO32-], [B(OH) 4-], [H3SiO4-], [HS-], [OH-], [H+] • Whenever there are significant amounts of any of these other species, they must be considered • Carbonate alkalinity = [HCO3-] + 2[CO32-] + [OH-] – [H+] • Directly convertible to [HCO3-] when it is >> than others • Measured by titration of solution with strong acid

  22. Total DIC = 10-1 M

  23. Alkalinity Titration • Determine end-point pH: • The pH at which the rate of change of pH per added volume of acid is at a maximum • Typically in the range 4.3-4.9 • Function of ionic strength • Reported as mg/L CaCO3 • HCO3- = alkalinity 0.82

  24. Determining Alkalinity by Titration Initial pH = 8.26 Rapid pH change Rapid pH change Slow pH change: Buffered Determine maximum pH change by: ΔpH ÷ mL acid added

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