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WATER QUALITY IN STREAMS AND RIVERS IS THE END PRODUCT OF ALL PROCESSES IN THE BASIN

WATER QUALITY IN STREAMS AND RIVERS IS THE END PRODUCT OF ALL PROCESSES IN THE BASIN. WATERSHEDS ARE THE KIDNEYS OF AN ECOSYSTEM. KIDNEY ANALOGOUS TO A WATERSHED. NITRATE EXAMPLE. Fingerprint water Isotopes Geochemical content Nutrients. Stoichiometry.

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WATER QUALITY IN STREAMS AND RIVERS IS THE END PRODUCT OF ALL PROCESSES IN THE BASIN

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  1. WATER QUALITY IN STREAMS AND RIVERS IS THE END PRODUCT OF ALL PROCESSES IN THE BASIN

  2. WATERSHEDS ARE THE KIDNEYS OF AN ECOSYSTEM

  3. KIDNEY ANALOGOUS TO A WATERSHED

  4. NITRATE EXAMPLE

  5. Fingerprint water Isotopes Geochemical content Nutrients

  6. Stoichiometry Stoichiometry is the accounting, or math, behind chemistry. Given enough information, one can use stoichiometry to calculate masses, moles, and percents within a chemical equation. AgNO3(aq) + NaCl(aq) ---> AgCl (s) + NaNO3(aq) reactants products

  7. LITHOSPHERE • Linkage between the atmosphere and the crust • Igneous rocks + acid volatiles = sedimentary rocks + salty oceans (eq 4.1)

  8. IMPORTANCE OF ROCK WEATHERING [1] Bioavailability of nutrients that have no gaseous form: • P, Ca, K, Fe • Forms the basis of biological diversity, soil fertility, and agricultural productivity • The quality and quantity of lifeforms and food is dependent on these nutrients

  9. IMPORTANCE OF ROCK WEATHERING [2] Buffering of aquatic systems -Maintains pH levels -regulates availability of Al, Fe, PO4 Example: human blood. -pH highly buffered -similar to oceans

  10. IMPORTANCE OF ROCK WEATHERING [3] Forms soil [4] Regulates Earths climate [5] Makes beach sand!

  11. NATURAL ACIDS • Produced from C, N, and S gases in the atmosphere • H2CO3 Carbonic Acid • HNO3 Nitric Acid • H2SO4 Sulfuric Acid • HCl Hydrochloric Acid

  12. CARBONIC ACID Carbonic acid is produced in rainwater by Reaction of the water with carbon dioxide Gas in the atmosphere.

  13. CARBONATE (DISSOLUTION) All of the mineral is completely Dissolved by the water. Congruent weathering.

  14. DEHYDRATION Removal of water from a mineral.

  15. OXIDATION Reaction of minerals with oxidation. An ion in the mineral is oxidized.

  16. HYDROLYSIS H+ replaces an ion in the mineral. Generally incongruent weathering.

  17. HYDROLYSIS • Silicate rock + acid + water = base cations + alkalinity + clay + reactive silicate (SiO2)

  18. HYDROLYSIS • Base cations are • Ca2+, Mg2+, Na+, K+ • Alkalinity = HCO3- • Clay = kaolinite (Al2Si2O5(OH)4) • Si = H4SiO4; no charge, dimer, trimer

  19. Mineral Solubility • Solubility - relative capability of being dissolved • Salt dissolution - solids break down in solution to yield ions • Example: Barium chloride BaCl2 BaCl2 (s) = Ba2+ + 2 Cl–

  20. Define K using the Law of Mass Action (“activity” in brackets): • Inside the [] are the measured concentrations • Multiply [] by number of atoms

  21. Solubility constant Ksp • Because the activity of the solid is 1, the equation becomes Ksp = [Ba2+]· [Cl–]2 • The equilibrium constant for the dissolution reaction is called the solubility product constant or Ksp.

  22. Measurements of Disequilibrium • It can be important to know whether a solution is saturated or undersaturated with respect to a mineral • Consider: AaBb = aA + bB • At equilibrium: Ksp = [A]a [B]b • How do we know the solution is in equilibrium with the mineral? Measure [A] and [B] in solution (activity product or ion activity product) and compare to Ksp

  23. Degree of saturation W • where [A] and [B] are for the solution, • which may or may not be in equilibrium with the mineral • W > 1 Supersaturated • W = 1 Saturated • W < 1 Undersaturated

  24. Problem:What is the degree of saturation of anhydrite in College Station tap water? • (Ca2+) = 3 mg/L = 0.003 g·L-1/40 g Ca·mol-1 = 0.000075 M • (SO42-)= 10 mg/L =0.010 g·L-1/96 g SO42-·mol-1 = 0.00010 M • T = 25°C • Assume ideal behavior (g = 1) • Write the reaction in terms of dissolution and make use of Ksp values CaSO4 = Ca2+ + SO42-

  25. We calculate the ion activity product in solution: IAP = [Ca2+][SO42-] = 0.000075 · 0.00010 = 7.5 x 10–9 = 10–8.1 • Degree of saturation Water is undersaturated with respect to annhydrite

  26. Calcite dissolution: • CaCO3 = Ca2+ + CO32– Is water undersaturated or oversaturated with respect to calcite? Get stalagmites/stalagtites? Or dissolve them? Tea pots: where does mineral deposits come from?

  27. But ions don’t behave ideally . . . • Concentration related to activity using the activity coefficient g, where [z] = gz (z) • The value of g depends on: • Concentration of ions and charge in the solution • Charge of the ion • Diameter of the ion • Ionic strength I = concentration of ions and charge in solution I = 1/2 Smizi2 • where mi = concentration of each ion in moles per kg, zi = charge of ion

  28. Activity and Concentration • Activity – “effective concentration” • Ion-ionand ion-H2O interactions (hydration shell) cause number of ions available to react chemically ("free" ions) to be less than the number present • Concentration can be related to activity using the activity coefficient g, where [z] = gz (z) • Activity coefficient gz 1 as concentrations  0 and tend to be <1 except for brines

  29. Carbonate Chemistry

  30. The Carbonate System • pH of most natural waters controlled by reactions involving the carbonate system • Groundwater and seawater chemistries are often poised near calcite equilibrium, with pH buffered by calcite dissolution and precipitation • Applications • Fate of CO2 from fossil fuels and other CO2 sources on the atmosphere • Effect of acid rain on lakes • Effect of acid mine drainage on rivers

  31. Carbonate System • Carbonate species are necessary for all biological systems • Aquatic photosynthesis is affected by the presence of dissolved carbonate species. • Neutralization of strong acids and bases • Effects chemistry of many reactions • Effects global carbon dioxide content

  32. PCO2 = 10–3.5 yields pH = 5.66 • What is 10–3.5? 316 ppm CO2 • What is today’s PCO2? ~368 ppm = 10-3.43 • pH = 5.63

  33. pH of Global Precipitation

  34. DIPROTIC ACID SYSTEM • Carbonic Acid (H2CO3) • Can donate two protons (a weak acid) • Bicarbonate (HCO3-) • Can donate or accept one proton (can be either an acid or a base • Carbonate (CO32-) • Can accept two protons (a base)

  35. TOTAL CARBONATE SPECIES (CT)

  36. OPEN SYSTEM • Water is in equilibrium with the partial pressure of CO2 in the atmosphere • Useful for chemistry of lakes, etc • Carbonate equilibrium reactions are thus appropriate

  37. Activity of Carbonate Species versus pH

  38. CARBONATE SPECIES AND pH

  39. Carbonate Buffering: Humans

  40. We can describe the formation and dissociation of carbonic acid through the following chemical and equilibrium equations

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