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

12. Fingerprint water Isotopes Geochemical content Nutrients

13. Rock Weathering

14. LITHOSPHERE Linkage between the atmosphere and the crust. Weathering results in: Igneous rocks + acid volatiles = sedimentary rocks + salty oceans

15. 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

16. IMPORTANCE OF ROCK WEATHERING [2] Buffering of aquatic systems -Maintains pH levels -regulates availability of Al, Fe, PO 4 Example: human blood. -pH highly buffered -similar to oceans

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

18. NATURAL ACIDS that WEATHER ROCK Produced from C, N, and S gases in the atmosphere H 2 CO 3 Carbonic Acid HNO 3 Nitric Acid H 2 SO 4 Sulfuric Acid HClHydrochloric Acid

19. 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. Keep track of atoms, molecules, and charge Calcite dissolution CaCO 3 + CO 2 + H 2 O  Ca 2+ + 2 HCO 3 - reactants products

20. TYPES OF CHEMICAL WEATHERING Carbonate weathering Dehydration Oxidation Hydrolysis

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

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

23. DEHYDRATION Removal of water from a mineral.

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

25. OXIDATION: REDOX REACTIONS Loss of electrons = Oxidation Gain of electrons = Reduction

26. For example, rusting * Oxidation of elemental iron to iron(III) oxide by oxygen 4 Fe + 3 O 2 → 2 Fe 2 O 3 Occurs in nature as the mineral hematite, and is the principal ore for iron

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

28. HYDROLYSIS Silicate rock + acid + water = base cations + alkalinity + clay + reactive silicate (SiO 2 )

29. HYDROLYSIS Base cations are –Ca 2+, Mg 2+, Na +, K + Alkalinity = HCO 3 - Clay = kaolinite (Al 2 Si 2 O 5 (OH) 4 ) Si = H 4 SiO 4 ; no charge, dimer, trimer

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

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

32. Solubility constant K sp –Because the activity of the solid is 1, the equation becomes K sp = [Ba 2+ ] · [Cl – ] 2 –The equilibrium constant for the dissolution reaction is called the solubility product constant or K sp.

33. Measurements of Disequilibrium It can be important to know whether a solution is saturated or undersaturated with respect to a mineral Consider: A a B b = aA + bB At equilibrium: K sp = [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 K sp

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

35. Problem: What is the degree of saturation of anhydrite in College Station tap water? (Ca 2+ ) = 3 mg/L = 0.003 g·L -1 /40 g Ca·mol -1 = 0.000075 M ( SO 4 2- ) = 10 mg/L = 0.010 g·L -1 /96 g SO 4 2- ·mol -1 = 0.00010 M T = 25°C Assume ideal behavior (  = 1) Write the reaction in terms of dissolution and make use of K sp values CaSO 4 = Ca 2+ + SO 4 2-

36. We calculate the ion activity product in solution: IAP = [Ca 2+ ][SO 4 2- ] = 0.000075 · 0.00010 = 7.5 x 10 –9 = 10 –8.1 Degree of saturation Water is undersaturated with respect to annhydrite

37. Calcite dissolution: CaCO 3 = Ca 2+ + CO 3 2– Is water undersaturated or oversaturated with respect to calcite? Get stalagmites/stalagtites? Or dissolve them? Tea pots: where does mineral deposits come from?

38. But ions don’t behave ideally... Concentration related to activity using the activity coefficient , where[z] =  z (z) The value of  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  m i z i 2 –where m i = concentration of each ion in moles per kg, z i = charge of ion

39. Activity and Concentration Activity – “effective concentration” Ion-ion and ion-H 2 O 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 , where[z] =  z (z) Activity coefficient  z  1 as concentrations  0 and tend to be <1 except for brines

40. Carbonate Chemistry

41. 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 CO 2 from fossil fuels and other CO 2 sources on the atmosphere –Effect of acid rain on lakes –Effect of acid mine drainage on rivers

42. 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

43. P CO2 = 10 –3.5 yields pH = 5.66 »What is 10 –3.5 ? 316 ppm CO 2 What is today’s P CO2 ? ~368 ppm = 10 -3.43 »pH = 5.63

44. pH of Global Precipitation

45. DIPROTIC ACID SYSTEM Carbonic Acid (H 2 CO 3 ) –Can donate two protons (a weak acid) Bicarbonate (HCO 3 - ) –Can donate or accept one proton (can be either an acid or a base Carbonate (CO 3 2- ) –Can accept two protons (a base)

46. TOTAL CARBONATE SPECIES (C T )

47. OPEN SYSTEM Water is in equilibrium with the partial pressure of CO 2 in the atmosphere Useful for chemistry of lakes, etc Carbonate equilibrium reactions are thus appropriate

48. First the CO2 dissolves according to: (1) CO 2 (g) ⇔ CO 2 (l) According to Henry’s Law, solubility increases as water temperature decreases

49. (2) CO 2 (l) + H 2 O (l) ⇔ H 2 CO 3 (l) Equilibrium is established between the dissolved CO 2 and H 2 CO 3, carbonic acid.

50. (3) H 2 CO 3 + H 2 O  H 3 O + + HCO 3 - Carbonic acid is a weak acid that dissociates in two steps. pK a1 (25  C) = 6.37

51. (4) HCO 3 - + H 2 O  H 3 O + + CO 3 2- pK a2 (25  C) = 10.25

53. Activity of Carbonate Species versus pH

54. CARBONATE SPECIES AND pH

55. Carbonate Buffering: Humans

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

57. Carbonic acid forms when CO 2 dissolves in and reacts with water: CO 2(g) + H 2 O = H 2 CO 3 »Most dissolved CO 2 occurs as “aqueous CO 2 ” rather than H 2 CO 3, but we write it as carbonic acid for convenience »The equilibrium constant for the reaction is: »Note we have a gas in the reaction and use partial » pressure rather than activity

58. »First dissociation: H 2 CO 3 = HCO 3 – + H +

59. FIRST REACTION

60. »Second dissociation: HCO 3 – = CO 3 2– + H +

61. SECOND REACTION

62. Variables and Reactions Involved in Understanding the Carbonate System

63. ALKALINITY refers to water's ability, or inability, to neutralize acids. The terms alkalinity and total alkalinity are often used to define the same thing.

64. Total alkalinity - sum of the bases in equivalents that are titratable with strong acid (the ability of a solution to neutralize strong acids) Bases which can neutralize acids in natural waters: HCO 3 –, CO 3 2–, B(OH) 4 –, H 3 SiO 4 –, HS –, organic acids (e.g., acetate CH 3 COO –, formate HCOO – )

65. Carbonate alkalinity Alkalinity ≈ (HCO 3 – ) + 2(CO 3 2– ) Reason is that in most natural waters, ionized silicic acid and organic acids are present in only small concentrations If pH around 7, then –Alkalinity ≈ HCO 3 –

66. Alk = OH – + HCO 3 – + 2CO 3 –2 + B(OH) 4 - - H + Bicarbonate dominates alkalinity of sea water

67. Gran Titration for Acid Neutralizing Capacity (total alkalinity) This method determines ANC by titration with 0.1 N Hydrochloric Acid between the pH range of 4.5 and 3.5 at which the contributions of organic acids, carbonate and bicarbonate are neutralized. Explicitly accounts for most organic acids

68. Charge Balance Fundamental principle of solution chemistry is that solutions are electronically neutral Sum of positive charges must equal the sum of negative charges in any sample  + =  -  + >  -, there is an unmeasured anion Equivalents: moles/L x valence

69. ION Percent Difference Quality Control  + -  - /  + +  - Normalizes the charge balance NADP guidelines for allowable error Error should be random and equal zero If error always positive, means there is an unmeasured anion (negative charge)

70. Alkalinity is routinely measured in natural water samples. By measuring only two parameters, such as alkalinity and pH, the remaining parameters that define the carbonate chemistry of the solution (P CO2, [HCO 3 – ], [CO 3 2– ], [H 2 CO 3 ]) can be determined.

71. Calcium (Ca++) and Magnesium (Mg++) are primarily responsible for hardness. However, in most waters, alkalinity and hardness have similar values because the carbonates and bicarbonates responsible for total alkalinity are usually in the form of Calcium carbonate or Magnesium carbonate. However, waters with high total alkalinity are not always hard, since the carbonates can be in the form of Sodium or Potassium carbonate. Alkalinity and hardness.

72. CLOSED CARBONATE SYSTEM Carbon dioxide is not lost or gained to the atmosphere Total carbonate species (C T ) is constant regardless of the pH of the system Occurs when acid-base reactions much faster than gas dissolution reactions Equilibrium with atmosphere ignored

73. How does [CO 3 –2 ] respond to changes in Alk or DIC? C T = [H 2 CO 3 *] + [ HCO 3 – ] + [CO 3 –2 ] ~ [ HCO 3 – ] + [CO 3 –2 ] (an approximation) Alk = [OH – ] + [HCO 3 – ] + 2[CO 3 –2 ] + [B(OH) 4 - ] – [H + ] ~ [HCO 3 – ] + 2[CO 3 –2 ] (a.k.a. “carbonate alkalinity”) So (roughly): [CO 3 –2 ] ~ Alk – C T C T ↑, [CO 3 –2 ] ↓ Alk ↑, [CO 3 –2 ] ↑

74. Diurnal changes in DO and pH What’s up?

75. Photosynthesis is the biochemical process in which plants and algae harness the energy of sunlight to produce food. Photosynthesis of aquatic plants and algae in the water occurs when sunlight acts on the chlorophyll in the plants. Here is the general equation: 6 H20 + 6 CO2 + light energy —> C6H12O6 + 6 O2 Note that photosynthesis consumes dissolved CO2 and produces dissolved oxygen (DO). we can see that a decrease in dissolved CO2 results in a lower concentration of carbonic acid (H2CO3), according to: CO2 + H20 H2CO3 (carbonic acid) As the concentration of H2CO3 decreases so does the concentration of H+, and thus the pH increases.

76. Cellular Respiration Cellular respiration is the process in which organisms, including plants, convert the chemical bonds of energy-rich molecules such as glucose into energy usable for life processes. The equation for the oxidation of glucose is: C6H12O6 + 6 O2 —> 6 H20 + 6 CO2 + energy As CO2 increases, so does H+, and pH decreases. Cellular respiration occurs in plants and algae during the day and night, whereas photosynthesis occurs only during daylight.

77. Rock Cycle

82. Carbonate weathering Hydrolysis

89. %A R. M. Garrels %A F. T. MacKenzie %T Origin of the chemical composition of some springs and lakes %B Equilibrium Concepts in Natural Water Systems %E R. G. Gould %S Am. Chem. Soc. Adv. Chem. Ser. %V 67 %P 222-242 %D 1967 Classic reference on geochemical weathering