1 / 24

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.

brita
Download Presentation

Dissociation and pH

An Image/Link below is provided (as is) to download presentation Download Policy: Content on the Website is provided to you AS IS for your information and personal use and may not be sold / licensed / shared on other websites without getting consent from its author. Content is provided to you AS IS for your information and personal use only. Download presentation by click this link. While downloading, if for some reason you are not able to download a presentation, the publisher may have deleted the file from their server. During download, if you can't get a presentation, the file might be deleted by the publisher.

E N D

Presentation Transcript


  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

More Related