1. Streamwater Chemistry 1) Dissolved major ions 2) Dissolved nutrients 3) Suspended and dissolved organic matter 4) Dissolved gases 5) pH

2. 1) Dissolved major ions TDS (Total dissolved solids) = sum of all dissolved major ions [TABLE 4.2] –Regional variation TDS of 120 mg/L is world average[= 0.12 grams per 1000 grams water = 0.12 ppt] TDS of <500 mg/L is drinking water standard for U.S. Lower Colorado: TDS > 800 mg/L[0.8 ppt] How to measure TDS? –Evaporation –Conductivity – electrical conductance of water due to dissolved ions.

3. Salinity -- all anions and cations (effectively same as TDS in freshwater)FW < 0.5 ppt Ocean ave. = 35 ppt Hardness = sum of Ca 2+ and Mg 2+ –In U.S. used as synonym for alkalinity Alkalinity - Quantity of compounds that shift pH > 7. HCO 3 -, CO 3 2-, OH - units: mg/l (ppm) CaCO 3, or meq/l Alkalinity often used as surrogate for stream fertility … production of crustaceans (gammarids, crayfish) often higher with high Ca 2+ concentrations; aquatic insect production less sensitive. Fig. 4.6 shows salmonid production across a alkalinity gradient. - low alkalinity limits production - variation at high alkalinity What other environmental factors correlated with high alkalinity?

4. Sources of TDS? –Why is streamwater much more concentrated than rainwater? [FIG. 4.1] –Contact with minerals in soil Weathering of rock –carbonate rocks (limestone) »high in Ca 2+, Mg 2+, HCO 3 - –igneous or non-carbonate rocks (granite, slate, sand) Groundwater inputs (long contact with soil/rock)

5. Biological Effects –Fig. 4.7 shows effects of road salt on water salinity (Cl - ) along rural-urban gradient. Note these concentrations are just for Cl - ; total TDS would be much higher.

6. 2) Dissolved nutrients (N,P,C,Si) N P Fig. 13.3: N and P vary with land use type Forest to agriculture gradient Human inputs NO 3 in fertilizer PO 4 in animal wastes NO 3 in acid rainwater Human alteration of N and P export to oceans globally

7. 3) S uspended and dissolved organic matter Seston - suspended particulate matter, including plankton, organic detritus, and inorganic material. Dissolved Organic Material (DOM) – material < 0.45 micrometers, includes leachates from living organisms and soils, and decaying detritus

8. 4) dissolved gases (N 2, O 2, CO 2 ) and pH N 2 –source for N-fixing cyanophytes (blue-green algae) –gas bubble disease in fish where N 2 supersaturated (e.g., below large reservoirs) Analogous to “the bends” O 2, CO 2 –Factors controlling concentration Solubility temperature [TABLE 4.1] atmospheric pressure (altitude) –Supersaturation? turbulent mixing biological activity –photosynthesis Concentration in atmosphere at sea level O 2 = 21% CO 2 = 0.03%

9. What is pH and what controls it? pH of various liquids, rain, and lakes Natural gradients in pH Fig. 4.8: Sampling across streams in acidic regions of southern England microarthropods macroarthropods Main point? Fewer species adapted to low pH. What is it? –Negative log 10 of [H + ] If [H + ] = 4.5 × 10 −4 mol/L, pH = 3.35. Why is it important? –Life tolerance (~4.5 - ~9.5)

10. Biological effects of excessive [H + ] Examples (from text): Invertebrates: species composition changes along pH gradient in Swiss streams Fish: Brook trout decline while blacknose dace and sculpin can be eliminated by low pH in northeastern US Loss of body Na + and failure to acquire Ca 2+ Damage to respiratory surfaces (fish gills, mayfly gills) and egg development Leaching of toxic Aluminum from soils into streams (Fig. 4.9)

11. Sources of Acidity –Natural Acidification Poorly buffered soils (non-calcareous soils) [see equations] Humic substances (dissolved organic material from wetlands, etc.) –Anthropogenic acidification Addition of NO 3 -, SO 4 2- in “acid rain” What makes streams “vulnerable” to acidification?

12. Why is “acid rain” << pH 5.6? –Strong acids disassociate in water H 2 SO 4 2H + + SO 4 2- HNO 3 H + + NO 3 - - Carbonic acid forms and dissociates into weak acid. How much more acidic than neutral? [10 7 / 10 5.6 = 10 1.4 ] = 25 #2:H 2 O + CO 2 H 2 CO 3 HCO 3 - + H + But most streams are not acidic! Why is pH of “pure” rain only 5.6? What is the pH of distilled water? #1:H 2 O H + + OH - pH = 7 = - log 10 [10 -7 ]

13. What makes water acidic? addition of H + –what are sources? Acids (e.g., carbonic acid (Eqn #2)) Addition of CO 2 –from atmosphere, groundwater What makes water more “alkaline” (pH > 7)? addition of OH - –what are sources? Reactions of water with bicarbonate (HCO 3 - ) and carbonate (CO 3 2- ) ions –Where does carbonate and bicarbonate come from??? #5:CO 3 2- + H 2 O HCO 3 - + OH - #4:HCO 3 - + H 2 O H 2 CO 3 + OH - #2:H 2 O + CO 2 H 2 CO 3 HCO 3 - + H + from watershed!! (groundwater contact with limestone, (CaCO 3 ) a source of HCO 3 - )

14. The form of dissolved inorganic C depends on pH. –CO 2 free carbon dioxide –HCO 3 - bicarbonate –CO 3 2- carbonate Most streams in pH range of 6.5-9, and HCO 3 - dominates. If H + added to stream, neutralized by OH - formed from reaction of water with HCO 3 - (Eqn #3) or with CO 3 2- (Eqn #4) and pH does not change much. Adding enough H + can “use up” OH - provided by CO 3 2- or HCO 3 - and lower pH, eventually producing dissolved CO 2. #2:H 2 O + CO 2 H 2 CO 3 HCO 3 - + H +

15. How does Acid Rain + Stream Water = no change in pH? –OH - neutralizes H + and more OH - forms immediately from reaction of CO 3 2- or HCO 3 - with water! –pH will not change until supply of CO 3 2- or HCO 3 - exhausted Why are some streams more susceptible? –Limestone geology (CaCO 3 is source of HCO 3 - ) –More acidic rainfall (humans) most streams Bicarbonate Buffering System  Streams with high alkalinity (HCO 3 - or CO 3 2- ) can “hold” much H + without notable change in pH. #3:H 2 O + CO 2 H 2 CO 3 HCO 3 - + H + CO 3 2- + H +

16. A carbonate(d) twist (Eqn #6) Carbonic acid (from rain) reacts with limestone in soil: H 2 CO 3 + CaCO 3 Ca 2+ + 2HCO 3 - Calcium ion reacts with abundant HCO 3 - in stream to form Calcium bicarbonate: –CaCO 3 can precipitate out of stream water … under what conditions? –Removing CO 2 drives the equation to the right. –How can CO 2 be removed?? #6: Ca 2+ + 2HCO 3 - Ca(HCO 3 ) 2 CaCO 3 + H 2 O + CO 2

17. In “hard” waters: CO 2 removed in two ways: 1) Biological activity –Shoreline algal photosynthesis (mostly lakes) Chara (skunk weed) –Removes CO 2 and becomes encrusted with CaCO 3 –What happens at night? #6: Ca 2+ + 2HCO 3 - Ca(HCO 3 ) 2 CaCO 3 + H 2 O + CO 2

18. 2) Physical-chemical processes where “excess” dissolved CO 2 vented e.g., Travertine terraces (e.g., Mammoth Hot Springs, Yellowstone) #6: Ca 2+ + 2HCO 3 - Ca(HCO 3 ) 2 CaCO 3 + H 2 O + CO 2 What would happen in stream below a dam if water rich in calcium bicarbonate were released from the hypolimnion of a deep reservoir during summer stratification? Supersaturated CO 2 in subterranean water degasses upon contact with atmosphere Hint: Deep reservoir water is supersaturated with CO 2

19. Synopsis: CO 2 dissolves into surface water to equilibrium HCO 3 - and CO 3 2- enter through surface/ground water Controls on pH? 1) buffering reactions of carbonic acid 2) amount of carbonate and bicarbonate derived from rock weathering (produces OH - ) 3) buffering reactions also influenced by salinity, temperature, but we’re not concerned with that here Bicarbonate Buffering System deceptively simple Wetzel: “… in alkaline, hard water lakes, often twice the content of Ca 2+ and HCO 3 - found than predicted on the basis of chemical equilibria.”