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Water Science

Water Science. Electroneutrality pH, Alkalinity, Acidity Reversible Reactions / Equilibrium Equations Common-Ion & Secondary-Salt Effect Carbonate Equilibrium System. Electroneutrality. For most waters, including potable ∑ cation charges = ∑ anion charges

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Water Science

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  1. Water Science • Electroneutrality • pH, Alkalinity, Acidity • Reversible Reactions / Equilibrium Equations • Common-Ion & Secondary-Salt Effect • Carbonate Equilibrium System

  2. Electroneutrality • For most waters, including potable • ∑ cation charges = ∑ anion charges • ∑ [C] x chargec = ∑ [A] x chargeA e.g., pure water, [H+] = [OH-] e.g., rain water, [H+] = [OH-] + [HCO3-] + 2[CO3-2] [ ] = molar concentration = moles / liter (one mole = 6.02x1023) C = cation & A = anion chargeC = number of charges per ion C chargeA = number of charges per ion A

  3. pH • Way of reporting H+ concentration • pH = - log {H+} • {} = concentration in terms of activity, i.e., the effective concentration • pH = - log [H+]for dilute solutions • [] = Molar concentration (Moles/Liter) • [H+] = 10-pH • pH = -log[H+]

  4. Why do we care about pH? • Biological systems • Extremes are disruptive • Corrosivity • Equilibrium relationships involving H+

  5. Freshwater... Marine ... Drinking... Soda? 6.5 - 9... Good for aquatic life 6.5 - 8.5... Good for aquatic life 5 – 9 ~ 2 - 4 Good pH ranges

  6. Alkalinity • Capacity of water sample to neutralize an... • Capacity of a water sample to take H+ without significant change in... • Common ions that give water alkalinity • HCO3-, CO32- & OH- acid pH

  7. Acidity • Capacity of water sample to take OH- without significant change in pH • Opposite of Alkalinity • Capacity of water sample to neutralize a base • Common ion – H+ • For Acid Mine Drainage • Acidity =~f(Fe2+, Fe3+, Al, Mn, and H+)

  8. pH, Alkalinity and Acidity • A solution with a given pH can have different values of Alkalinity or Acidity • Because chemicals may be present that can react with or release H+ or OH- • The lower the alkalinity (or acidity), the easier it is to change pH

  9. Reversible Reactions • aA + bB ↔ cC + dD • A and B can react to form C and D • C and D can react to form A and B • After sufficient time has passed, equilibrium is reached • Equilibrium can be perturbed by adding more reactant or product

  10. Reversible Reactions In Solution • At equilibrium • Reactants become products • Products become reactants • At equal rate For example: HCO3- ↔ H+ + CO32-

  11. Reversible Reactions: Gas - Solution • At equilibrium • Gaseous chemical dissolves • Dissolved chemical volatilizes • At equal rate Gas Dissolved For example: CO2g ↔ CO2aq

  12. Reversible Reactions: Solution - Precipitate • At equilibrium • Precipitate is formed • Precipitate is dissolved • At equal rates For example: Ca2+ + CO32- ↔ CaCO3

  13. For a generic reversible reaction aA + bB ↔ cC + dD Equilibrium Equation K is constant at given temperature & relatively low concentration

  14. What is water? ... Chemical equation?... Equilibrium equation?... At 25 C, Kw= ... Even if other sources of H+ (acids) or OH- (bases) are present Which substance predominates?... H2O, H+ and OH- H2O ↔ H+ + OH- [H+ ] [OH-] = Kw Leave out [H2O], as it is nearly constant, 55.56 M 10-14 H2O Example - Pure Water

  15. Follow these steps Write down species... Write electroneutrality eq... Write equilibrium equation.. Solve equations... H2O, H+ and OH- [H+] = [OH-] [H+ ] [OH-] = 10-14 M2 Sub [OH-] = [H+] [H+]2 = 10-14 M2 [H+] = 10-7 M pH = - log [H+] ... = - log (10-7) = 7 pH of Pure Water @ 25 C?

  16. Now, add other chemicals • Add source of OH- • Systems finds new equilbrium, with less H+ and more OH- • Add source of H+ • Systems finds new equilbrium, with more H+ and less OH- • Either way, [H+] [OH-] = Kw still holds

  17. Precipitation of Limestone • [CaCO3-2] ↔ Ca+2 + CO3-2 • [Ca+2] [CO3-2] = Ksp • Ksp = Solubility product • Do not include precipitate in equilibrium equation • [Ca+2] [CO3-2] < Ksp? • [Ca+2] [CO3-2] > Ksp? CaCO3 dissolves… CaCO3 precipitates

  18. Carbonate, Alkalinity, Acid Rain • Carbonate solids are a common source of alkalinity in natural waters (as they dissolve) • Lakes and stream with alkalinity can accept acid rain without big pH changes • the alkalinity buffers the acid • Acid rain can kill water bodies that don’t have alkalinity

  19. Solubilities • In general • Metals are less soluble at higher pH • Metal hydroxides are less soluble than Metal carbonates which are less soluble than Metal bicarbonates • There are exceptions! • Aluminum compounds tends to be soluble at high and low pH • CaOH more soluble than CaCO3

  20. Common-Ion & Secondary-Salt Effect • Common-Ion Effect • Add common ion to solution, reduce solubility • e.g., [Ca+2] [CO3-2] = Ksp. Add Na2CO3. Ca and CO3 will precipitate as CaCO3 • Secondary-Salt Effect • Add other ions to solution, increase solubility • E.g., Add NaCl to increase solubility of CaCO3

  21. Carbonate Equilibrium System (CES) • Related to • pH, alkalinity, hardness, buffering, water and wastewater treatment, acid rain, Acid mine drainage... • Keeps natural waters in pH range good for living creatures • Species • Carbon dioxide, CO2 • Carbonic acid, H2CO3 • Bicarbonate, HCO3- • Carbonate, CO3-2 • Natural systems: Solids containing Bicarbonate, Carbonate, Hydroxide, and (though rarely) certain Oxides (CaO),…

  22. Sources • CO2 from atmosphere • Byproduct of • biological or chemical reactions • combustion • Solids containing carbonate • e.g., limestone • Byproduct of sea life

  23. Overview Phase Reversible Reactions Gas in atmosphere CO2 g Gas-Liquid (Henry’s Law) CO2 aq H2CO3 In solution Dissolved in liquid HCO3 CO3 Precipitation Solid in liquid Solid CO3,…

  24. CO2 g ↔ CO2 aq CO2 aq + H2O ↔ H2CO3 H2CO3↔ H+ + HCO3- HCO3- ↔ H+ + CO3-2 Solid Source, e.g., limestone Ca+2 + CO3-2↔ CaCO3 Henry’s Law Reversible Reaction Reversible Reaction Reversible Reaction Precipitation Reaction CES Reactions

  25. CES Equilibrium Equations (at 25oC) • CO2 aq = CO2 g / 1637 atm • Where CO2 g is in atm & CO2 aq is in mole fraction • [H2CO3] / [CO2 aq] = 1.58 x 10-3 • [H+] [HCO3-] / [H2CO3*] = 4.47 x 10-7 M • [H+] [CO3-2] / [HCO3-] = 4.68 x 10-11 M • [Ca+2] [CO3-2] = 10-8.42 M2 • Where [] means Molar concentration, mol/l

  26. Predominate Species? < 4.5: CO2 4.5 - 8.3: HCO3- 10-11: CO3-2 >11: OH- 100 HCO3 CO3 mg/L as CaCO3 50 CO2 OH 6.5 8.5 11 Carbonate System and pH Total alk. = 100 mg/L as CaCO3 pH A Typical Natural Water

  27. Carbonate Systems • Open or Closed • Open - in equilibrium with the atmosphere • Requires large contact surface • Closed - not in equilibrium • With or Without CO3 solids • With - CO3 solid in contact with solution • Gives water ability to buffer, i.e., resist pH change • Without - no CO3 solids in contact

  28. CES – pH of Rain • Carbonate equilibrium system • Open • W/O solid source of CO3 • Use Electroneutrality & Equilibrium Equations to estimate pH

  29. pH of rainwater • pH of rainwater is controlled by CO2 in Atmosphere • Nitrogen - 0.781 atm • Oxygen - 0.209 atm • Argon - 0.0093 atm • CO2 - 0.00033 atm • Misc. - 0.0004 atm • TOTAL - 1.0 atm Let’s work example…

  30. CO2 g CO2 aq H2CO3 H+ HCO3- H+ CO3- Natural pH of Rain • CO2 in atmosphere naturally “shifts” pH of rain from 7 to ~5.6, we can estimate this Lowers pH “Acid rain” is rain with pH below 5.7, i.e., caused by human activity, e.g., relasing SOx compounds by burning coal

  31. Science Summary • Electroneutrality • Sum Cations = Sum Anions • pH, Alkalinity, Acidity • pH = - log [H+] • Alkalinity: HCO3-, CO32- & OH- • Acidity: H+ • Equilibrium Reactions

  32. Equilibrium Equations (25oC) CO2 g = 1637 atm • CO2 aq [H2CO3] / [CO2 aq] = 1.58 x 10-3 [H+] [HCO3-] / [H2CO3*] = 4.47 x 10-7 M (Where [H2CO3*] = [H2CO3 aq] + [CO2 aq], used to make measurement easier) Don’t need for this problem [H+] [CO3-2] / [HCO3-] = 4.68 x 10-11 M [Ca+2] [CO3-2] = 10-8.42 M2 No solid source

  33. Solution: pH of Rain (1) • Atmosphere to rain - Henry’s Law:CO2 aq = CO2 g / 1637 atm • CO2 aq = 0.0003 atm / 1637 atm • CO2 aq = 1.83x10-7 (mole fraction) • Convert to M: [CO2 aq] = CO2 aq x Mw • Mw = molar density of water • [CO2 aq] = 1.83x10-7 x 55.56 mol/l • [CO2 aq] = 1.02x10-5 mol/l = 1.02x10-5 M

  34. Solution: pH of Rain (2) • Reaction with water forms carbonic acid - Reversible Reaction (Equilibrium Eq.): [H2CO3 aq] = 1.58x10-3 M • [CO2 aq] • [H2CO3 aq] = 1.58x10-3• 1.02x10-5 M • [H2CO3 aq] = 1.61x10-8 M • [H2CO3* aq] = [CO2 aq] + [H2CO2 aq] • [H2CO3* aq] = 1.02x10-5 M + 1.61x10-8 M • [H2CO3* aq] = 1.02 x 10-5 M

  35. Solution: pH of Rain (3) • Dissociation of carbonic acid - Reversible Reaction (Equilibrium Eq.):[H+] [HCO3-] = 4.47 x 10-7 M • [H2CO3*] • Two variables? Use electroneutrality: [H+] = [OH-] + [HCO3-] + 2[CO3-2] • Rain water is acidic, so [OH-] & [CO3-2] will be small (see next slide), giving [H+] ≈ [HCO3-], leading to • [H+]2 = 4.47x10-7 M • [H2CO3*] • [H+] = (4.47x10-7 M • 1.02x10-5)0.5 = 2.13x10-6 M

  36. Aside: Log Concentration Diagram Closed System w/o solid CaCO3 Which species predominate at pH 4.5? 6.5? 8? 12? Note: total concentration as carbon CT = [C] = 0.001 (pK1 & pK2 are equilibrium equation constants given in the “p-notation”, like pH)

  37. Solution: pH of Rain (4) • pH = - log [H+] = -log (2.13x10-6) = 5.67 • Check assumptions? Use equilibrium equations for water and carbonate: • [OH-] = 10-14 M2 / [H+] = 10-14 M2 / 2.13x10-6 M = 4.68x10-9 M (small, OK)) • [CO3-2] = 4.68x10-11 M • [H+] / [HCO3-] = 4.68x10-11 M • 2.13x10-6 M / 2.13x10-6 M = 4.68x10-11 M (small, OK)

  38. pH of Rain • CO2 in atmosphere goes up? • pH of rain drops • Use this method for different conditions? • Need to check assumptions about relative concentration of [OH-] & [CO3-2]

  39. CO2, pH & Oceans? • Oceans can be CO2 sink or source • Currently: sink, taking ~ 30% of anthropocentric CO2 • Primary mechanisms • Carbonate equilibrium system (Henry’s Law,…) • Biological Pump • Living creatures take up carbon, some are trapped in sea bed • Too much CO2? • pH drop of oceans could effect sea organisms • From 1751 to 1994 surface ocean pH estimated to have dropped from ~ 8.18 to 8.10

  40. Ocean pH Change • ~half Anthropogenic CO2 has been absorbed by ocean so far • Currently absorbing ~1/3 • 22 M tons / day • Could drop to 7.6 • Drastic effect on shell-forming organisms National Geographic (2014) “Ocean Acidification”, ocean.nationalgeographic.com.

  41. Red Oak Mine Site

  42. Side View pH = 4.4 Zero Alkalinity AMD Seep

  43. Plan View Not to scale

  44. Red Oak Seep seep discharge

  45. Red Oak Remediation

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