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
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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
LITHOSPHERE • Linkage between the atmosphere and the crust • Igneous rocks + acid volatiles = sedimentary rocks + salty oceans (eq 4.1)
IMPORTANCE OF ROCK WEATHERING  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
IMPORTANCE OF ROCK WEATHERING  Buffering of aquatic systems -Maintains pH levels -regulates availability of Al, Fe, PO4 Example: human blood. -pH highly buffered -similar to oceans
IMPORTANCE OF ROCK WEATHERING  Forms soil  Regulates Earths climate  Makes beach sand!
NATURAL ACIDS • Produced from C, N, and S gases in the atmosphere • H2CO3 Carbonic Acid • HNO3 Nitric Acid • H2SO4 Sulfuric Acid • HCl Hydrochloric Acid
CARBONIC ACID Carbonic acid is produced in rainwater by Reaction of the water with carbon dioxide Gas in the atmosphere.
CARBONATE (DISSOLUTION) All of the mineral is completely Dissolved by the water. Congruent weathering.
DEHYDRATION Removal of water from a mineral.
OXIDATION Reaction of minerals with oxidation. An ion in the mineral is oxidized.
HYDROLYSIS H+ replaces an ion in the mineral. Generally incongruent weathering.
HYDROLYSIS • Silicate rock + acid + water = base cations + alkalinity + clay + reactive silicate (SiO2)
HYDROLYSIS • Base cations are • Ca2+, Mg2+, Na+, K+ • Alkalinity = HCO3- • Clay = kaolinite (Al2Si2O5(OH)4) • Si = H4SiO4; no charge, dimer, trimer
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–
Define K using the Law of Mass Action (“activity” in brackets): • Inside the  are the measured concentrations • Multiply  by number of atoms
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.
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
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
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-
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
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?
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
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
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
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
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
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)
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
We can describe the formation and dissociation of carbonic acid through the following chemical and equilibrium equations