Unit 1: Module 2/3 Part 5 - Major Ions and Nutrients January 2004 LAKE ECOLOGY Unit 1: Module 2/3 Part 5 - Major Ions and Nutrients January 2004

Modules 2/3 overview Goal – Provide a practical introduction to limnology Time required – Two weeks of lecture (6 lectures) and 2 laboratories Extensions – Additional material could be used to expand to 3 weeks. We realize that there are far more slides than can possibly be used in two weeks and some topics are covered in more depth than others. Teachers are expected to view them all and use what best suits their purposes. Goal: Provide a practical introduction to lake ecology. This is not a comprehensive limnology course. Rather, it is a “crash course” to be integrated with other tool-oriented WOW modules for initial training in technical areas related to water resource management. Lecture time: 2-3 weeks of classroom instruction with weekly lab/field experience. Slides: Divided into 6 subtopics. Note – Subtopics 4-6 use WOW data and visualization tools first introduced in Subtopic 4 – the density stratification discussion. Also, the module introduces lake biota before discussing physical and chemical data, although some instructors may want to reverse this order. Status (January 2004) Lecture – WOW staff review in progress; a few more slides are in prep; some graphic design needed Lab – in prep; focus will be on “traditional” field surveys comparing local ponds/lakes to each other and to WOW lakes.

Modules 2/3 outline Introduction Major groups of organisms; metabolism Basins and morphometry Spatial and temporal variability – basic physical and chemical patchiness (habitats) Major ions and nutrients Management – eutrophication and water quality A 2 week-module can only highlight the basics of limnology. Students should be referred to the variety of introductory and advanced limnology texts now available. Some of these include: Cole, G.A. 1994. Textbook of limnology. 4th edition. Dodds, W.K. 2002. Freshwater Ecology: Concepts and environmental applications. Academic Press, San Diego, CA. USA. Horne, A. J. and C.R. Goldman. 1994. Limnology. 2nd Edition. McGraw-Hill, Inc. New York. Hutchinson, G.E. 1957. 4 volumes . A treatise on limnology. John Wiley & Sons, New York. Mason, C.F. 1996. Biology of freshwater pollution. 3rd edition. Longman House Publ., Essex, UK. McComas, S. 1993. Lake Smarts: The first lake maintenance handbook. Terrene Institute, Washington, D.C., USA. Monson, B.A. 1992. A primer on limnology (2nd edition). Water Resources Center, University of Minnesota, St. Paul, MN, USA. Moss, , B., J. Madgwick and G. Phillips. 1996. A guide to the restoration of nutrient-enriched shallow lakes. W.W.Hawes, UK. NALMS. 2002. Lake and reservoir guidance manual. North American Lake Management Society, Madison , WI (http://nalms.org) Schmitz, R.J. 1996. An introduction to water pollution biology. Gulf Publ. Co., Houston, TX, USA. Welch, E.B. 1992. Ecological effects of wastewater: Applied limnology and pollutant effects. 2nd edition. Chapman & Hall, London, UK. Wetzel, R.G. 2002. Limnology 3rd Edition. Academic Press, San Diego, CA. Wetzel, R.G. and G.E. Likens. 2002. Limnological analyses. 3rd edition. Springer-Verlag, NY,NY, USA. ELAINE: I ACCIDENTALLY ZAPPED WHAT RICH HAD AS BULLETS IN ORIGINAL SLIDE. PROBLEMS WITH HIS SLIDE – BOX DIDN’T FILL PAGE – USE MASTER SLIDE DEFAULT AND IT WILL. CAPITALIZED SOME STUFF THAT SHOULDN’T HAVE BEEN ITALICS UNNECESSARY (TEXT IN ITALICS ALSO UNNECESSARY IN MY OPINION).

5. Water chemistry: Gases, major ions & nutrients

5. Water chemistry: Gases, major ions & nutrients Oxygen (O2) Carbon dioxide (CO2) Nitrogen (N2) Hydrogen sulfide (H2S) Major ions (anions and cations) Nutrients (phosphorus and nitrogen) Additional information is contained in the materials prepared for Lab Modules 2/3 and 4/5

Water chemistry: gases What are the ecologically most important gases ? O2 CO2 N2 H2S

Gas solubility The maximum amount of gas that can be dissolved in water (100% saturation) is determined by temperature, dissolved ion concentration, and elevation solubility decreases with temperature “warm beer goes flat” solubility decreases with higher dissolved ion content (TDS, EC25, salinity) “DO saturation is lower in saltwater than freshwater (for the same temperature, solids “drive out” gases) Why does elevation affect the concentration of dissolved gases? The temperature effect is much greater than that due to salinity. The range of salinities between most natural “freshwater” has relatively little effect

Water chemistry: O2 ~ 21% of air Very soluble (DO) Highly reactive and concentration is dynamic Involved in metabolic energy transfers (PPr & Rn) Major regulator of metabolism (oxic-anoxic) Aerobes (fish) vs anaerobes (no-fish, no zoops) Types of fish Salmonids = high DO (also coldwater because of DO) Sunfish, carp, catfish = low DO (also warmwater) PPr = primary productivity Rn = respiration

O2 variability Diel (24 hr) variation due to ____________? Seasonal variation due to _____________ ? Diel (24 hr) variation due to day to night and sunny to cloudy changes Seasonal variation due to : temperature, productivity, and daily insolation (sunlight) differences

Major sources of O2 Sources Photosynthesis (phytoplankton, periphyton, macrophytes) Air from wind mixing Inflows tributaries may have higher or lower DO groundwater may have higher or lower DO Diffusion (epilimnion to hypolimnion and vice versa) Estimating Primary Productivity (PPr) and Respiration (Rn) Careful measurements of DO over periods of an hour over the course of a day can be used to estimate both PPr and Rn in lakes and streams (with some assumptions about diffusive exchange of O2 with the atmosphere). Another method involves comparing changes in DO over periods of ~ an hour or more in bottles of water exposed to ambient sunlight and temperature as compared to bottles incubated in darkness. These techniques are routinely taught in limnology and stream ecology field and lab courses and are beyond the scope of this course. However, the respiration measurement is essentially the same as the measurement of BOD (biochemical oxygen demand that estimates the organic matter concentration) although a “seed” of wastewater or freeze-dried bacteria are usually added when BOD is being measured (see Module 9 for details and references)

Major sinks of O2 Sinks Respiration bacteria, plants, animals; water and sediments Diffusion to sediment respiration Outflow (tributary or groundwater)

Gases: wind mixing from storms Oxygen from a storm – How many mixing “events” can you find for Halsteds Bay in Lake Minnetonka, MN in this 1 year record? Compare Halsteds Bay to the West Upper site in Lake Minnetonka using both the DxT and Color Mapper tools. For additional information on this lake go to http://www.lakeaccess.org . Our RUSS data and Hennepin Park's water chemistry data from summer 1999 showed that sudden mixing events at Halsteds Bay could have dramatic effects on water quality. Severe thunderstorms with high winds and tornado warnings ripped through the Twin Cities area in August and September 1999, leading to sudden mixing and re-aeration events at the Halsteds Bay RUSS site in Lake Minnetonka. The event had profound water quality consequences by disrupting stable thermal stratification and introducing high nutrient/low oxygen bottom water into the upper sunlit layer of the bay resulting in an obnoxious algae bloom some days later. This, in turn, led to further changes in dissolved oxygen, pH and turbidity. The nearby West Upper Lake water column responded differently to the wind event and produced much smaller, if any, serious water quality effects. In fact, there are economically important land and lake management implications that derive from these data in terms of how to most efficiently restore the degraded water quality in Halsteds Bay. Restorative management decisions involving millions of dollars of tax monies are currently being discussed for Halsteds Bay. Such storm-related events were suspected, but prior to our RUSS data, were not documented due to their transient nature and the danger of manual sampling during severe weather. In fact, several of these events were observed in August 1999 (see figures below) and again in August and September 2000. In each case, temperature profiles didn't tell us the state of mixing of the bay since variations from surface to bottom were only about 1oC. However, dissolved oxygen (and to a lesser extent pH and EC25) clearly indicated that the lower half (~5 m) of the water column dramatically changed from extreme anoxia (no oxygen) to >75% saturation and then back again to anoxia over intervals of ~ 24 hours in mid and late August/early September. During the initial mixing, the influx of anoxic water to the "epilimnion" actually decreased the level of DO to ~5 mg/L, which could potentially harm fish communities. In addition, water samples collected on August 25, about 10 days after the first mixing event, showed that about 3200 kg (over 7000 lbs of P) was suddenly injected into the upper sunlit euphotic zone. This represented an areal phosphorus load of about triple the annual load estimated to enter Halsteds Bay from Six Mile Creek, its major tributary. The sudden input of P appears to have then caused an increase in algal growth, seen as a chlorophyll increase over the same manual monitoring interval (Barten and Vlach 1999, L. Minnetonka Annual Monitoring Report). Recall that 7000 lbs of P can potentially lead to over 3,500,000 lbs of algae ! These new data, that would likely not have been acquired without remote sensing, suggest that water quality in nutrient enriched lakes of intermediate depth (perhaps 5-10 meters) may often be controlled by weather events. The data also indicate that watershed Best Management Practices (BMPs), alone, may not be successful in improving the water quality of these bays without concurrently reducing internal, in-lake P-loading. Of course this does not diminish the importance of preventing increases in external P-loading from old and new developments over the long-term. However, it does offer important insights into the cost-effective management of water quality for 6 bays on the west and north end of the lake that have shown significant downward trends in water quality over the past 5 years.

Gases: seasonal wind mixing Oxygen varies seasonally and the entire water column lake may be fully saturated at certain times. How often did this happen in Ice Lake, MN in this 5+ year record? Complete Mixing at Ice Lake, MN Spring: No-98,99,01 Yes- 00,02,03 Fall: No-00 (?? – Difficult to say) Yes- 98,99,01,02 Also Students should also examine the data more closely for each period using the DxT tool and the color mapper or profile plotter. DO should compared to temperature, pH and EC25. Students should determine what the data for each parameter suggests in regard to mixing. Which are the most sensitive and why?

O2: Human significance Not a direct threat to humans Directly affects fish physiology and habitat Indirectly affects fish and other organisms via toxicants associated with anoxia: H2S NH4+ (converts to NH4OH and NH3 above ~pH 9) Indirectly affects domestic water supply H2S (taste and odor) Solubilizes Fe (staining) Indirectly affects reservoir turbines Via H2S corrosion and pitting (even stainless steel) Via regulation of P-release from sediments (mediated via Fe(OH)3 adsorption) Diel (24 hr) variation due to day to night and sunny to cloudy changes diel variations are clearly visible in the Halsteds Bay, Lake Independence, Medicine Lake and Lake Onondaga data sets using the Color Mapper/Profile Plotter from ~ June through August. These are all eutrophic – hypereutrophic systems Seasonal variation due to: temperature, productivity, and daily insolation (sunlight) differences Use the DxT tool to view annual or multi-year data sets for any of the WOW lakes

Gases: N2 ~ 78% of air Concentrations in water usually saturated because it is nearly inert Supersaturation (>100 %) can occur in reservoir tailwaters from high turbulence May be toxic to fish (they get “the bends) N2 -fixing bacteria and cyanobacteria (blue-green “algae”) convert it to bio-available NH4+ Denitrifying heterotrophic bacteria convert NO3- to N2 and/or N2O under anoxic conditions See N-cycle set of slides later in this Module

Gases: CO2 Only about 0.035% of air (~ 350 ppm) Concentration in H2O higher than expected based on low atmospheric partial pressure because of its high solubility Gas (at 10oC) Concentration @ 1 atm (mg/L) Concentration @ normal pressure (mg/L) N2 23.3 18.2 O2 55.0 11.3 CO2 2319 0.81 The table shows : CO2 is much more soluble than O2 which is more than twice as soluble as N2. N2 is still higher than O2 or CO2 at normal pressures because its concentration (i.e. its partial pressure) in air is so much higher. How much fizz is in a bottle of soda pop? Do an experiment by sealing a 2 gallon baggie over the top of a 1 pint bottle with a rubber band. Shake it and hold on tight. How long does your soda pop fizz after shaking it?

CO2 reactions in water <1% is hydrated to form carbonic acid: CO2 + H2O H2CO3 Some of the carbonic acid dissociates into bicarbonate and hydrogen ions which lowers the pH: H2CO3 HCO-3 + H + As the pH rises, bicarbonate increases to 100% at a pH of 8.3. Above this, it declines by dissociating into carbonate: HCO-3 CO3-2 + H+ These interconversions assume that free hydroxide (OH-) concentrations are negligible, as they usually are unless there is an industrial effluent

Inorganic - C equilibria H2CO3 HCO3 CO3 pH Fraction of carbon species Source of generic diagram: www.swbic.org/education/env-engr/carbonate/carbonate.html Note – 100% CO2 for pH< ~ 4.5; 100% bicarbonate for pH ~ 8 and 100% carbonate for pH > ~12

Inorganic - C: Major sources and sinks Atmospheric CO2 (invasion) Respiration and other aerobic and anaerobic decomposition pathways in the water and sediments Groundwater from soil decomposition products Groundwater from volcanic seeps Sinks: pH dependent conversions to bicarbonate and carbonate Precipitation of CaCO3 and MgCO3 at high pH Photosynthesis Relevance to WOW RUSS data: Water strata with high rates of photosynthesis are characterized by increased pH since CO2 removal rates are high The hypolimnion in a stratified lake accumulates CO2 from the decomposition of settling organic matter (from respiratory and anaerobic metabolism). This causes pH to decrease and EC25 to increase due to the accumulation of bicarbonate at the typical pH values of 6-8 found in most hypolimnia Look for these types of variations in the summer daily profiles (using the color mapper/profile plotter) for the more productive WOW lakes and over seasonal time frames for all of the lakes. Very careful measurements of pH and CO2 changes over the course of a day have been used to estimate both net photosynthesis and respiration in lakes and streams. The method is difficult, especially in better buffered systems (references and further details may be found in most limnology texts).

CO2 supersaturation – killer Lake Nyos In 1986, a tremendous explosion of CO2 from Lake Nyos, in Cameroon, West Africa, killed >1700 people and livestock up to 25 km away. Dissolved CO2 seeps from volcanic springs beneath the lake and is trapped in deep water by hydrostatic pressure. Nearby Lake Manoun is similar in nature Although unconfirmed, a landslide probably triggered the gas release Images and text courtesy of Dr. George Kling at the University of Michigan. His website (http://www.biology.lsa.umich.edu/~gwk/research/nyos.html) describes the phenomenon in great deal and he has done extensive research on these lakes. Lower image courtesy of Dr. Michel Halbwachs (http://perso.wanadoo.fr/mhalb/nyos/index.htm for detailed information). Permission granted 10-21-03 Visit http://www.biology.lsa.umich.edu/~gwk/research/nyos.html and http://perso.wanadoo.fr/mhalb/nyos/index.htm for detailed information

Soda pop chemistry www.saddleback.cc.ca.us/faculty/thuntley/ms20/seawaterprops2/sld013.htm Graphic from Dr. Tony Huntley at Saddleback College, Mission Viejo, CA (http://www.saddleback.cc.ca.us/faculty/thuntley/ms20/seawaterprops2/sld013.htm) Permission requested on Oct 20, 03 Another schematic of the inorganic carbon reactions as they pertain to the ocean (identical to freshwater although the system is much better buffered than most freshwaters.

CO2 and the inorganic carbon system Carbon dioxide diffuses from the atmosphere into water bodies and can then be incorporated into plant and animal tissue It is also recycled within the water with some being tied up in sediments and some ultimately diffusing back into the atmosphere Fixed carbon also enter the water as “allocthonous” particulate and dissolved material See Stream and Lake Primers for additional information on food webs and carbon cycling. For details refer to Limnology and Stream Ecology texts listed in the LINKS section.

CO2 and the inorganic carbon system - 2 Alkalinity, acid neutralizing capacity (ANC), acidity, carbon dioxide (CO2), pH, total inorganic carbon, and hardness are all related and are part of the inorganic carbon complex

CO2 chemistry: Alkalinity Alkalinity – the ability of water to neutralize acid; a measure of buffering capacity or acid neutralizing capacity (ANC) Total Alkalinity (AlkT) = [HCO3-] + 2[CO32-] +[OH-] - [H+] Typically measured by titration with a strong acid. The units are in mg CaCO3/L for reasons relevant to drinking water treatment (details in Module 9) Can be used to estimate the DIC (dissolved inorganic carbon) concentration if the [OH-] Conversely, direct measurements of DIC by infrared analysis or gas chromatography, together with pH and the carbon fractionation schematic can be used to estimate alkalinity (* see slide notes) Notes- The interest in alkalinity from the perspective of drinking water treatment plants is that it is important to maintain high enough alkalinity and calcium concentrations in the water system to develop a moderate amount of calcium carbonate precipitation on the inside of pipes. This can reduce the leaching of harmful lead, copper and zinc. Also, water with pH below ~ 8 has increasing amounts of carbonic acid which is corrosive to pipes and other infrastructure. Other considerations include having too much alkalinity (and calcium) that creates excess deposits of calcium carbonate (the mineral “marl”) in home fixtures and excessive hardness (see next slide) that interferes with soap sudsing. DIC is typically measured by acidifying a water sample in a closed chamber to convert all the carbonate and bicarbonate to carbon dioxide, and then sparging it out of solution with a stream of helium or argon gas gas and sweeping it through a column to remove water vapor and then to a detector in either an infrared analyzer or a gas chromatograph. Individual fractions can then be estimated using previous inorganic-C equilibrium graph or by calculation using actual equilibrium constants (“pK” values).

Alkalinity and water treatment Advanced wastewater treatment (domestic sewage) Phosphorus nutrient removal by adding lime (Ca(OH) 2) or calcium carbonate (CaCO3) As pH increases >9, marl precipitates adsorbed PO4-3 Settle and filter the effluent to obtain 90-95% removal Used for particle (TSS) removal also Drinking water treatment For TSS removal prior to disinfection Acid-rain mitigation to whole lakes Lime or limestone added as powdered slurry to increase impacted lake pH Also broadcast aerially to alkalize entire watersheds The interest in alkalinity from the perspective of drinking water treatment plants is that it is important to maintain high enough alkalinity and calcium concentrations in the water system to develop a moderate amount of calcium carbonate precipitation on the inside of pipes. This can reduce the leaching of harmful lead, copper and zinc. Also, water with pH below ~ 8 has increasing amounts of carbonic acid which is corrosive to pipes and other infrastructure. Other considerations include having too much alkalinity (and calcium) that creates excess deposits of calcium carbonate (the mineral “marl”) in home fixtures and excessive hardness (see next slide) that interferes with soap sudsing.

CO2 chemistry: Hardness Hardness - the total concentration of multi-valent (i.e. >2) cations Ca+2 + Mg+2 + Fe +3 (when oxic) + Mn+2 (when oxic); all other multivalent cations are typically considered to be negligible Sources- Minerals such as limestone (Ca and Mg) and gypsum (Ca) Water softeners and other water treatment processes such as reverse osmosis and ion exchange Evaporation can increase hardness concentration Drinking water effects (no real health effects) Soap scums and water spots on glasses and tableware Deposits (scaling) can cause clogging problems in pipes, boilers and cooling towers See specific methodology for estimating hardness in Module 9 (Lake Lab methods)

Water chemistry – Major ions SiO2 < 1 Table Reference - these values are one of several literature estimates of the chemical composition of the world’s rivers (from Cole, G.A. 1979. Textbook of Limnology (2nd edition). C.V. Mosby Company, St. Louis, MO, USA. General Lake Chemistry In the absence of any living organisms, a lake would contain an array of molecules and ions from the weathering of soils in the watershed, the atmosphere, and the lake bottom. The chemical composition of a lake is therefore, fundamentally a function of its climate and its basin geology. Each lake has an ion balance of the 3 major anions and 4 major cations (silica, actually in the ionized form of silicate is usually a relatively minor fraction of the anions). Lakes with high concentrations of the ions, Ca2+ and Mg2+ are called hardwater lakes, while those with low concentrations of these ions are softwater lakes. Concentrations of other ions, especially bicarbonate, are highly correlated with the concentrations of the hardness ions, especially Ca2+. The ionic concentrations influence the lake’s ability to assimilate pollutants (e.g., acidification) and maintain nutrients in solution. For example, calcium carbonate (CaCO3) in the form known as marl, can precipitate phosphate from the water, thereby removing this important nutrient from the water. Of the anions, bicarbonate typical dominates at moderate pH levels (~ 7-9) and calcium is usually the major cation. Note: plant nutrients such as nitrate, ammonium and phosphate that can cause algae and weed overgrowth usually occur at 10’s or 100’s of parts-per-billion and along with other essential micronutrients usually represent <1% of the actual amount of cations or anions present in the water

Major ion concentrations - freshwater Anions mg/L Cations HCO3- 58.4 Ca+2 15.0 SO4-2 11.2 Mg+2 4.1 Cl- 7.8 Na+ 6.3 SiO2 13 K+ 2.3 NO3- ~1.0 Fe+3 ~0.7 Total = ~91.4 anions + ~28.4 cations = ~ 120 mg/L (TDS) Table Reference - these values are from Livingstone 1963 as reported in Cole, G.A. 1979. Textbook of Limnology (2nd edition). C.V. Mosby Company, St. Louis, MO, USA (Table 12-1).

Nutrients – phosphorus Essential for plant growth Usually the most limiting nutrient in lakes Derives from phosphatic rock – abiotic, unlike nitrogen No gas phase, but can come from atmosphere as fugitive dust Adsorbs to soils Naturally immobile unless soil is eroded or excess fertilizer is applied Phosphorus moves with sediments The most limiting nutrient is one that can limit the growth of an algal cell. Other factors such as light, temperature, or the maximum growth rate can also limit the rate of growth. This is explained in more detail in slide 37.

Nutrients – phosphorus Not toxic Algae have physical adaptations to acquire phosphorus High affinity (low k) APA Storage Luxury uptake Single redox state Phosphorus cycle is closely linked to the iron (Fe) cycle High affinity (low k): This refers to the fact that some plants, (algae included) are better than others at taking up nutrients that are in short supply in the water. Having high “affinity” means that they are very good at this. “Low k” refers to having a low half-saturation constant for this uptake – the process being often described by a beast called the Michaelis-Menton equation. Since this is far above the introductory level of this module, the interested student should find a microbiology or microbial ecology or limnology text (such as those listed in the Notes for the introductory slides for this module). APA: Alkaline phosphatase activity. Under conditions of nutrient limitation, algae can increase the rate of alkaline phosphatase activity, allowing algae to use organophosphate or inorganic polyphosphates as alternative sources of phosphorus. This gives these algae a competitive advantage under low nutrient conditions. Storage: Some algae are capable of storing during periods when ortho-PO4 exceeds growth requirements. Blue-green algae store the P internally as polymeric polyphosphate bodies. Bloom species may load up on phosphate during periods that don’t favor growth (i.e. nighttime) then return to surface waters that are low in ambient phosphate but high in light. Luxury uptake: the accumulation of a nonlimiting resource above the levels required to maintain the current growth rate. Redox: phosphorus ions exist only as PO4–3 Significance of P cycle being linked to Fe cycle: phosphate concentration is strongly affected by iron redox reations.

Phosphorus – basic properties No redox or respiration reactions directly involved (organisms are not generating energy from P chemistry) PO4–3 highly adsorptive to cationic sites (Al+3, Fe+3, Ca+2) Concentration strongly affected by iron redox reactions Ferric (+3) – insoluble floc Ferrous (+2) – soluble, unless it reacts with sulfide, causing FeS to precipitate See nomenclature PDF file in Unit 1- Modules 9 and 11 (Lake and Stream Lab Chemistry) regarding various fractions of limnologically relevant phosphorus.

Phosphorus levels in the environment Major factors affecting phosphorus levels, cycling, and impacts on water quality include: Soil properties Land use and disturbance Transport associated with runoff

Where does phosphorus come from? Although some phosphorus may derive from atmospheric deposition of dust and soil particles containing adsorbed phosphate, the great majority usually comes from the watershed in close association with the transport of soil particles. Untreated or partially treated domestic or industrial wastewater effluents are also major sources and may be point (end-of-pipe) or nonpoint (diffuse). See http://www.lakeaccess.org for information about a watershed-scale lawn fertilizer phosphorus experiment conducted in the Medicine Lake watershed in the Minneapolis, MN metropolitan area. There are also links to other sources of information on this topic.

Phosphorus – external sources Nonpoint sources Watershed discharge from tributaries Atmospheric deposition Point sources Wastewater Industrial discharges

Phosphorus – nonpoint sources Watershed discharges from tributaries Strongly tied to erosion (land use management) Stormwater runoff (urban and rural) Agricultural and feedlot runoff On-site domestic sewage (failing septic systems) Sanitary sewer ex-filtration (leaky sewer lines) Atmospheric deposition Often an issue in more pristine areas Arises from dust, soil particles, waterfowl

Phosphorus – point sources Wastewater Municipal treated wastewater Combined sewer overflows (CSOs) Sanitary sewer overflows (SSOs) Industrial discharges

Phosphorus – internal sources Mixing from anoxic bottom waters with high phosphate levels is closely tied to iron redox reactions O2 > 1 mg/L – Insoluble ferric (+3) salts form that precipitate and settle out, adsorbing PO4-3 O2 < 1 mg/L (anoxic) – ferric ion reduced to soluble ferrous ion (Fe+2) – allowing sediment phosphate to diffuse up into the water Wind mixing (storms and fall de-stratification) can re-inject high P water to the surface, causing algal blooms Students should be referred to any limnology, microbial ecology, or geochemistry textbook for a more thorough discussion of the link between phosphorus and iron cycling. Not all lakes have enormous releases of phosphate from the sediment interstitial (pore) water during late summer or under-ice anoxia, because there may not be much iron in the system. For the purposes of this curriculum, i.e., training students for careers in applied lake management, it is fair to say that most stratified lakes will exhibit substantial increases in bottom water phosphate, which accumulates after the lakes become anoxic at the sediment-water interface. Note also that as redox values fall and the sediments become increasingly reduced, there will be large amounts of sulfide produced as hydrogen sulfide gas (H2S) from the energy metabolism of heterotrophic sulfate-reducing bacteria. If ferrous iron concentrations are relatively high, it will react with the sulfide to form ferrous sulfide – the black gooey muck you find in mudflats and marshes. This can also affect P-cycling in some cases, but turns out to be even more important in transformations involving heavy metals at polluted sites, where toxic metal reactions with sulfide act to bind the metals and reduce their exposure as toxicants to aquatic organisms. Texts will also point out that manganese (Mn) cycling can also influence P-cycling, much the same as iron does, because of the close similarity in redox chemistry between these two ions. They may often co-occur as well, although Mn levels are usually substantially lower than Fe. Where lakes have been historically polluted by high P-inputs, this reservoir of P can exceed annual watershed inputs (see Halsteds Bay and Shagawa data on WOW website for case studies) In summary: No redox or respiration reactions are directly involved with phosphorus chemistry. Organisms are not generating energy from P-transformations. PO4–3 ion is highly adsorptive to cationic sites (Al+3, Fe+3, Ca+2) on soil particles and in the sediments. P concentrations are strongly affected by Fe redox reactions: Ferric (+3) = insoluble floc; ferrous (+2) = soluble, unless it reacts with sulfide and FeS precipitates. Soil properties and transport associated with water movements are the major factors affecting phosphorus levels and cycling.

Phosphorus – Lake budget Note to instructor: Left-hand figure: Use this figure to highlight inputs and outputs of phosphorus. Right-hand figure: Discuss the relative magnitudes of the pools and sources and sinks for phosphorus. This particular cartoon shows a thermally stratified lake with increased P in the hypolimnion. It seems to show a large inflow of P relative to outflow, suggesting that a lot of phosphorus must be settling out of the epilimnion. However, if this were the case, one might expect the sediment phosphorus pool to be much larger than is indicated. Note: This is not necessarily a great schematic, but it does allow the instructor to discuss what might be going on in a real lake. It was taken from the Maine NEMO (Nonpoint [source] Education for Municipal Officials) Extension Program slideshow, but is a generic schematic of unknown original origin.

Nutrients – phosphorus cycle Major pools and sources of P in lakes “Natural” inputs are mostly associated with particles Wastewater is mostly dissolved phosphate P is rapidly removed from solution by algal-bacterial uptake or by adsorption to sediments Modified from Horne and Goldman. 1994. Limnology. McGraw Hill. Note to Instructor: Point out to students that phosphate is also called ortho-P, dissolved inorganic-P, soluble reactive phosphorus (SRP), etc. Step students though each “pool” in the figure to the right. Highlight that phosphate concentrations are usually very low in the euphotic zone due to rapid assimilation by algae and bacteria, even in relatively productive systems. High levels of phosphate in the upper water column usually indicate an influx of high P water, such as from a wastewater plant or agricultural drainage. It may also be a short-term result of transient mixing of high-P anoxic bottom waters in productive shallow lakes.

Phosphorus cycling – major sources Sewage Dissolved Tributaries and deposition Particulate Erosion Sediments Particulate and dissolved Note – These statements are all generalizations.

Phosphorus cycling – internal recycling Rapid PO4-3 recycling Bacterial uptake Algal uptake Adsorption to particles Detritus mineralization Zooplankton excretion Fish excretion The pool of phosphate in the oxygenated, sunlit portion of the water column is extremely dynamic. Typically, relatively low concentrations of dissolved orthophosphate are maintained by a dynamic balance between rapid uptake by P-starved algae and bacteria and regenerative processes such as the decomposition of dead organic material from plants and animals and excretion by animals (mostly zooplankton, but fish may also be important).

Phosphorus cycle – major transformations The whole phosphorus cycle Modified from Horne and Goldman, 1994. Limnology. McGraw Hill.

Nitrogen – basic properties Nitrogen is relatively scarce in some watersheds and therefore can be a limiting nutrient in aquatic systems Essential nutrient (e.g., amino acids, nucleic acids, proteins, chlorophyll) Differences from phosphorus Not geological in origin Unlike phosphorus, there are many oxidation states

Nitrogen – biologically available forms N2 – major source, but usable by only a few species Blue green algae (cyanobacteria) and anaerobic bacteria Nitrate (NO3-) and ammonium (NH4+) – major forms of “combined” nitrogen for plant uptake Also called dissolved inorganic nitrogen (DIN) Total nitrogen (TN) – includes: DIN + dissolved organic nitrogen (DON) + particulate nitrogen Dissolved organic nitrogen is typically mostly refractory (difficult to break down) material. Particulate nitrogen includes living and dead material, such as algal and bacterial cells.

Nitrogen – general properties Essential for plant growth Not typically limiting but can be in: Highly enriched lakes Pristine, unproductive lakes located in watersheds with nitrogen-poor soils Estuaries, open ocean Lots of input from the atmosphere Combustion NO2, fertilizer dust

Nitrogen – general properties Mobile – in the form of nitrate (soluble), it goes wherever water flows Ammonium (NH4+) adsorbs to soil particles Blue green algae can fix nitrogen (N2) from the atmosphere Nitrogen has many redox states and is involved in many bacterial transformations

Nitrogen – sources Atmospheric deposition Wet and dry deposition (NO3- and NH4+) Combustion gases (power plants, vehicle exhaust, acid rain), dust, fertilizers Streams and groundwater (mostly NO3-) Sewage and feedlots (NO3- and NH4+) Agricultural runoff (NO3- and NH4+) Regeneration from aquatic sediments and the hypoliminion (NH4+)

Nitrogen - toxicity Methemoglobinemia – “blue baby” syndrome > 10 mg/L NO3--N or > 1 mg/L NO2--N in well water Usually related to agricultural contamination of groundwater NO3- – possible cause of stomach/colon cancer Un-ionized NH4+ can be toxic to coldwater fish NH4OH and NH3 at high pH N2O and NOx – contribute to smog, haze, ozone layer depletion, acid rain Blue baby disease is of concern world-wide and in particular in agricultural areas. The cause has typically been attributed to residential drinking water wells having high nitrate-nitrogen concentrations caused by contamination in runoff from fields fertilized with nitrogen fertilizer. The disease is actually due to the nitrite (NO2-) ion binding to blood hemoglobin competitively with oxygen. This reduces the oxygen content of the blood and is life threatening. The blood actually appears brownish and the overall coloring of the child is bluish. Although nitrite is the ion that actually causes the problem, it is the more oxidized form of nitrogen, nitrate, that is usually the route of the problem. This is because the gut flora of young children, typically < 6 months old, have not yet become very acidic. Conditions are favorable when nitrate is in high concentrations (> 10 mg N/L) for anoxic, denitrifying bacteria to biologically reduce enough nitrate to nitrite to cause the symptoms of the disease. Note that the U.S. and World Health Organization drinking water criterion for nitrite is only 1 mg N/L. This is 10% of the nitrate standard. Cancer (suspected carcinogenicity) – A number of studies in the past decade have suggested a link between high levels of nitrate in drinking water and stomach and/or colon cancer, but to date there appears to be insufficient consensus for a change in the regulatory criterion, which has remained unchanged for many decades. Un-ionized ammonia (NH4OH + NH3) – At typical pH values in most natural waters (~ 6.5 – 8.5), ammonia is dissolved in water as the ammonium ion (NH4+). In this form, it is virtually non toxic unless extremely high values are reached (> 10s of mg N/L). As the pH increases, the equilibrium reaction moves to convert ammonium to ammonium hydroxide and then, ultimately, to free ammonia gas – both of which are very toxic to aquatic organisms. However, except when total ammonium levels are extremely high – many parts per million – the pH usually must be well over 9 or even higher before un-ionized ammonia levels become a problem. High temperature and high salinity (electrical conductivity) exacerbate the toxicity effect.

Nitrogen – many oxidation states Unlike P there are many oxidation states Organisms have evolved to make use of these oxidation-reduction states for energy metabolism and biosynthesis -3 + 1 + 2 + 3 + 5 NH4+ N2 N2O NO2 NO2- NO3-

Nitrogen – bacterial transformations Organic N NH4+-N Heterotrophic ammonification or mineralization. Associated with oxic or anoxic respiration. NH4+-N NO3- Involves oxygen (oxic). Autotrophic and chemosynthetic ("burn” NH4+-N to fix CO2). NO3- N2 (gas) Anoxic process. Heterotrophic. ("burn" organic matter and respire NO3-, not O2). N2 (gas) Organic N Some blue green algae are able to do this. Decomposition Nitrification Denitrification Nitrogen fixation

Nutrients – nitrogen cycle Nutrients- The Nitrogen Cycle The nitrogen cycle is really only a cycle in a global sense and not necessarily “in balance” in a given aquatic ecosystem. This type of schematic is useful for illustrating how external inputs of plant-available nitrogen, such as ammonium and nitrate from agricultural and lawn fertilizers, domestic and industrial wastewater, urban runoff, and atmospheric deposition are incorporated into lake biota and how they may ultimately be removed from the system. A detailed description of the various processes and transformations of nitrogen is beyond the scope of this curriculum. However, we have prepared a number of slides for use if instructors wish to go deeper than an introductory water science class. Nitrogen exists in a variety of oxidation states from strongly reduced (-3 when in the form of ammonium or the amine group) to strongly oxidized (+5 such as the nitrate-N). These various states are available to plants, animals and bacteria to varying degrees and are also used and transformed by micro-organisms in their energy metabolisms as well as in various biosynthetic reactions (such as protein and nucleic acid synthesis). Oxygen is a critical “switch” for many N-cycle processes Some of these reactions can only occur under strongly reducing conditions where there is absolutely no oxygen – such as the process of nitrogen fixation where molecular oxygen irreversibly denatures the enzyme nitrogenase that essentially acts to split the N2 molecule to form two ammonium ions. However, organisms have evolved a variety of “clever” structures, specialized cells (heterocysts in some blue-green algae/cyanobacteria), root nodules in alder trees, and flake formation in some cyanobacteria to maintain such an environment in a highly oxygenated epilimnion. Other processes such as denitrification are strictly anaerobic/anoxic but the organisms (heterotrophic bacteria) can switch from an oxygen based metabolism to a nitrate based metabolism when all the oxygen has been consumed – such as in the hypolimnion and a few centimeters into the sediments. Nitrification is an energy yielding process where bacteria essential “burn” or oxidize ammonium to nitrate to obtain energy that is used to “fix” dissolved carbon dioxide into organic matter – chemosynthesis in contrast to plant photosynthesis that uses light energy instead of chemical energy for reducing the carbon. Mineralization is just another word for the process of organic matter being converted to inorganic matter – decomposition. modified from Horne and Goldman. 1994. Limnology. McGraw Hill.

Chemical forms of nitrogen in aquatic systems organism-N + detrital-N + dissolved organic-N Org–N NO3- NO3- Dissolved inorganic-N (DIN) Fixed or available-N Nitrate: major runoff fraction Ammonium: basic unit for biosynthesis NH4 + NH4 + NO2- Nitrite: usually transient NO2- N2 = largest reservoir but cannot be used by most organisms Modified from the original with permission from Dr. Stefano Bernasconi, Geologisches Institut, ETH-Zentrum, CH-8092 Zuerich, Switzerland The scale at the bottom represents oxidation states for the nitrogen atom in the various molecules. -3: This is the most chemically reduced state in natural systems. Its inorganic form is gaseous ammonia, which readily dissolves in water to form the ammonium ion at circumneutral and low pH. It has the amine group (NH2-) as its most common form in organic matter. As pH increases, such as may occur from industrial discharges of caustic hydroxides or from very intense algal photosynthesis during the daytime, the ammonium can begin to form ammonium hydroxide and ultimately free ammonia gas – both of which are very toxic to aquatic animals. Together they are called un-ionized ammonia and are included as specific water quality criteria for lakes and streams. High total ammonium, high pH (typically values need to exceed 9 before it can become significant), and high temperature together can lead to problems. This is a common problem where fish densities are high, such as fish hatcheries, commercial aquaculture facilities, and home aquaria. Other cold water fish, such as trout and salmon, are the most sensitive fish and the basis of most regulations. In pristine unproductive systems, the pool of dissolved organic-N is typically the largest of any pool except for gaseous N2. This is because it is mostly comprised of refractory compounds that are difficult for bacteria to decompose (mineralize) to ammonium. The more easily (labile) broken down organic-N compounds, such as amino acids, nucleic acids and most proteins, are rapidly assimilated by bacteria or mineralized to ammonium. In these systems, the particulate fraction of organic nitrogen (and carbon as well, since all organic matter by definition contains carbon) usually represents about 10% of the dissolved fraction. 0: nitrogen gas is the by far the largest pool of N on the planet and usually also in water, which would have a concentration of ~ 18 mg N/L (100% saturation) at a temperature of 10 oC (50 oF). +1 and +2: Nitrous oxide (N2O, often called laughing gas) and nitric oxide are gases that are usually in low concentrations in water and are intermediates in various bacterial transformations. N2O is a potentially important greenhouse gas and so has received increased attention in recent years. NO2- is an important contributor to the formation of photochemical smog. +3: Nitirite (NO2-) also typically occurs at low concentrations as an intermediate in various bacterial N transformations. It is readily taken up and assimilated by bacteria and algae as a nutrient and can also be chemically oxidized to nitrate. It is quite toxic to fish and invertebrates at part per million levels and is a leading cause of tropical fish mortality in new home aquaria before the nitrifying bacteria community is well developed. +5: Nitrate (NO3-) is the predominant form of inorganic nitrogen found in most aquatic systems where the water has nmeasureable oxygen. N2 N2O NO2 gases -3 +5 +4 +3 +2 +1 -1 -2 Oxidation state

Functionally in the lab using filters… Total-N = particulate organic-N + dissolved organic-N + particulate inorganic-N + dissolved inorganic-N TN = PN + DON + DIN Dissolved inorganic-N = [Nitrate + Nitrite]-N + ammonium-N DIN = NO3-N + NO2-N + NH4-N Notes: Nitrate+nitrite are usually measured together. Nitrite is usually negligible. See nomenclature PDF file in Unit 1- Modules 9 and 11 (Lake and Stream Lab Chemistry) regarding various fractions of limnologically relevant nitrogen. TOTAL = PARTICULATE + DISSOLVED TOTAL = ORGANIC + INORGANIC Particulate inorganic-N: Usually negligible – there just aren’t nitrate or ammonium rocks, in general (fossil bat and bird guano are the exception, but are not widespread) Nitrite: Very reactive and usually present in low quantities except where wastewater inflows are present or in narrow strata in the hypolimnion near the oxic-anoxic boundary. It is rapidly taken up by algae and bacteria, may be oxidized to nitrate by certain nitrifying bacteria, and may also be chemically oxidized to nitrate even in a bottle on the shelf. Nitrate: Predominant form of inorganic nitrogen found in most aquatic systems where the water has measureable oxygen. Ammonium: The most chemically reduced form of inorganic-N in aquatic systems. It is a natural product of decomposition and excretion processes and is readily taken up and assimilated into biomass by algae, bacteria and macrophytes. It is the dominant fraction found in anaerobic (anoxic) zones and, if present in high concentrations in surface waters, is probably indicative of inflows polluted by domestic wastewater (sewage) or agricultural runoff (fertilizers).

Main N-cycle transformations Assimilation (algae + bacteria) Assimilation Denitrification Mineralization Org-N NO2- NO3- NH4+ Nitrification 1 (oxic bacteria) Nitrification 2 Anammox (anoxic bacteria) Ammonification Denitrification (anoxic bacteria) N2 - Fixation - Soil bacteria - Cyanobacteria - Industrial activity - Sulfur bacteria Modified from the original with permission from Dr. Stefano Bernasconi, Geologisches Institut, ETH-Zentrum, CH-8092 Zurich, Switzerland. Students are referred to a modern limnology, microbial ecology, or microbiology text for more in-depth information regarding the many microbial transformations of nitrogen. Because of its many oxidation states and its importance in biomolecules such as enzymes and nucleic acids, microorganisms have evolved a wide variety of synthetic and metabolic reactions that transform nitrogen between these states. The scale at the bottom represents oxidation states for the nitrogen atom in the various molecules. -3: This is the most chemically reduced state in natural systems. Its inorganic form is gaseous ammonia, which readily dissolves in water to form the ammonium ion at circumneutral and low pH and has the amine group (NH2-) as its most common form in organic matter. As pH increases, such as may occur from industrial discharges of caustic hydroxides or from very intense algal photosynthesis during the daytime, the ammonium can begin to form ammonium hydroxide and, ultimately, free ammonia gas – both of which are very toxic to aquatic animals. Together they are called un-ionized ammonia and are included as specific water quality criteria for lakes and streams. High total ammonium, high pH (typically values need to exceed 9 before it can become significant), and high temperature together can lead to problems. This is a common problem where fish densities are high, such as fish hatcheries, commercial aquaculture facilities, and home aquaria. Other cold water fish, such as trout and salmon, are the most sensitive fish and the basis of most regulations. In pristine, unproductive systems, the pool of dissolved organic-N is typically the largest of any pool, except for gaseous N2. This is because it is mostly comprised of refractory compounds that are difficult for bacteria to decompose (mineralize) to ammonium. The more easily broken down (labile) organic-N compounds, such as amino acids, nucleic acids, and most proteins, are rapidly assimilated by bacteria or mineralized to ammonium. In these systems, the particulate fraction of organic nitrogen (and carbon as well, since all organic matter by definition contains carbon) usually represents about 10% of the dissolved fraction. Zero: Nitrogen gas is the by far the largest pool of N on the planet and usually also the largest pool in water, which would have a concentration of ~18 mg N/L (100% saturation) at a temperature of 10 oC (50 oF). +1 and +2: Nitrous oxide (N2O, or laughing gas) and nitric oxide are gases that are usually in low concentrations in water and are intermediates in various bacterial transformations. N2O is a potentially important greenhouse gas and so has received increased attention in recent years. NO2 is an important contributor to the formation of photochemical smog. +3: Nitrite (NO2-) also typically occurs at low concentrations as an intermediate in various bacterial N transformations. It is readily taken up and assimilated by bacteria and algae as a nutrient and can also be chemically oxidized to nitrate. It is quite toxic to fish and invertebrates at part per million levels and is a leading cause of tropical fish mortality in new home aquaria before the nitrifying bacteria community is well developed. +5: Nitrate (NO3-) is the predominant form of inorganic N found in most aquatic systems where the water has measureable oxygen. gases N2 N2O NO2 -3 +5 +4 +3 +2 +1 -1 -2 Oxidation state

Whole lake N-budget N2 Algae Outflow NO3- NH4+ NO3- oxic anoxic NH4+ Tribs, GW, Precip DON, PON, NO3-, NH4+ N2 Ammonia volatilization N2-fixation Assimilation Outflow Algae NO3- NH4+ Denitrification DIN PON DON Nitrification Mineralization Mixing Sedimentation NO3- oxic anoxic Sedimentation NO2-, N2O NO Modified from the original with permission from Dr. Stefano Bernasconi, Geologisches Institut, ETH-Zentrum, CH-8092 Zurich, Switzerland. MAJOR SOURCES (INPUTS) = BLUE Tributaries Nitrate is usually the major fraction High ammonium suggests domestic or industrial wastewater, or agricultural drainage Urban stormwater runoff also contributes relatively high levels of both nitrate and ammonium. 2. Groundwater Usually nitrate High ammonium may indicate wastewater from onsite septic systems in addition to industrial wastes, agricultural runoff 3. Precipitation Nitrate and/or ammonium are the major fractions Major anthropogenic sources are combustion (NOx from power plants and vehicles), agricultural fertilizer (ammonium) Dry deposition can be as large a load as wet deposition (precipitation) for both nitrate and ammonium. Can be a major source of available N to relatively pristine lakes. 4. Mixing from the hypolimnion Partial mixing or complete turnover can re-introduce significant amounts of nitrate and/or ammonium into the euphotic zone when available dissolved inorganic-N (DIN) may be deficient. Ammonium will accumulate in anoxic strata, while under oxic conditions this mineralized ammonium will be largely nitrified to nitrate. Ammonium diffusion from anoxic sediments also contributes to higher concentrations in the deeper hypolimnion. MAJOR SINKS (OUTPUTS) = RED Tributary outflow of surface water Groundwater outflow (mixed layer concentrations may be modified by “filtration” through lake sediments, as water “exits” the lake Denitrification of nitrate to N2 gas in anoxic environments (deep hypolimnion and sediments) Permanent sedimentation of detrital-N to deep sediments Surficial sediments may be resuspended and mineralized to ammonium, but even in ultra-oligotrophic lakes there is at least some permanent deposition of organic material. Nitrification NH4+ Mineralization diffusion Surficial Sediments Burial Burial Deep Sediments

Nutrients – summer vertical profiles Oligotrophic Eutrophic NH4 T O2 NO3 NH4 T PO4 PO4 NO3 O2 Depth anoxia Important point: Oxygen concentration is a critical “switch” that regulates bacterial, plant, and animal metabolism, in turn controlling concentrations of the critical nutrients, nitrogen and phosphorus. SUMMER: Oligotrophic TEMPERATURE: Thermally stratified; well developed thermocline. O2 : Uniformly high in epilimnion; increases in metalimnion due to colder water, low demand, and some photosynthesis; relatively high throughout most of the hypolimnion; may be low or anoxic near the bottom. NO3–: Low in most of the water column because of low inputs and efficient uptake by nutrient-starved algae. Some accumulation deep in the water column, where it is too dark for algal uptake and also because of bacterial nitrification of ammonium from decomposing organic matter and excretion of ammonium from animals (O2 must be > ~1 mg/L for nitrification to occur). NH4+: Low in most of the water column because of low inputs and efficient uptake by nutrient-starved algae. Where there is some O2 but it is too dark for photosynthetic organisms to grow, some of the ammonium released from decomposition and animal excretion is nitrified by bacteria to nitrate. In the very deepest layer where O2 levels have dropped to zero, denitrifying bacteria can convert all of the pre-existing nitrate to N2 gas (and some N2O gas). PO4–: Phosphate also remains low throughout most of the water column due to low inputs coupled with efficient uptake by nutrient starved algae. When the bottom layer becomes anoxic, the oxidized “iron cap” on top of the sediments is chemically reduced and dissolves, allowing phosphate to diffuse up into the water from the sediments. Eutrophic: TEMPERATURE: Thermally stratified; well developed thermocline. Trophic state has almost nothing to do with the thermal regime of the lake. O2 : uniformly high in epilimnion; decreases dramatically to zero just below the thermocline. Entire hypolimnion is anoxic because of high rates of organic matter input that stimulates bacterial activity and respiration. NO3–: Likely to be relatively high in the mixed layer because of high rates of nitrate input (after all, high nutrient inputs are why it’s a eutrophic lake). However, it is rapidly and completely removed in anoxic water by bacterial denitrification to N2 and N2O. No nitrate can be converted from the large amount of ammonium that might build up, because anoxia prevents the growth of nitrifying bacteria (O2 must be > ~1 mg/L for nitrification to occur). NH4+: Likely to be relatively high in the mixed layer because of high rates of input from the watershed – again, high rates of available N and P are why it’s a eutrophic lake. The concentration builds up even higher below the thermocline and continues to increase steadily throughout the hypolimnion because – (1) the anoxia prevents it from being nitrified to nitrate, (2) there’s a lot of decomposing algae “raining” down from the productive, sunlit upper mixed layer; and (3) darkness and anoxia essentially shut down down any algal growth and uptake of nutrients. PO4–: Phosphate also is high in the mixed layer due to inputs from the watershed. It increases in the hypolimnion due to settling detritus from algal production in the upper water column. The bottom-most layer receives additional phosphate from the sediments that can diffuse out because there is no oxidized “iron cap.” Note: Not all lakes have enough iron in the water to exert total control over P-diffusion from the sediments. But most eutrophic lakes do respond this way. Medicine Lake, Lake Minnetonka (the deep West Upper station and the shallower Halsteds Bay) and Lake Independence in the Minneapolis, MN metropolitan area all behave like this for most of the ice-free season. Lake Onondaga, NY is another extremely productive lake with extremely high nutrient concentrations in the hypolimnion. All of these have near real-time data sets on the WOW website for temperature and oxygen (and also pH and specific electrical conductivity EC25). Water chemistry data is posted on WOW for the Minnesota Lakes (visit the DATA section). Lake Onondaga nutrient data can be found via links to Upstate Freshwater Institute reports at http://ourlake.org. Note: The interaction of iron oxidation state and phosphorus release from the sediments is extremely important in lake restoration and management (see also Module 24 on Lake Restoration). Lakes that have received excessive inputs of phosphorus for decades build up a large reservoir of phosphate in the sediments. Much of its release by diffusion into the overlying water is prevented because the phosphate adsorbs to “sticky,” oxidized iron complexes (mostly ferric hydroxide). Iron in the water column and surficial sediments that is exposed to oxygen quickly forms an insoluble floc that acts like a phosphorus trap. Even if millions of dollars are spent removing phosphorus from wastewater inflows and nonpoint sources, when the bottom water becomes anoxic in the summer, this enormous pool can essentially bleed large amounts of P into the water. It continues to build up throughout the stratified period. If the wind is strong enough, some upper hypolimnetic water, with its high concentration of phosphate, can be mixed up into the epilimnion. This combination of high P and high light can lead to a dramatic algal bloom. Usually, ammonium has also built up to high concentrations in the anoxic hypolimnion as well. Note: Studies of a number of nutrient polluted lakes that were targeted in the 1970s and 1980s for restoration by decreasing external nutrient inputs have now shown that this internal regeneration of P can persist for many decades. Shagawa Lake, one of the WOW lakes in Ely, MN is one such system (see the lake description in the DATA section of the website). Sediments in Halsteds Bay in Lake Minnetonka, MN can release the equivalent of more than an entire year of phosphorus loading from its major tributary in just a few mid-summer mixing events (see details at http://wow.nrri.umn.edu/wow/notes/Minnetonka_animation.html). Note: Oxidized iron formation can be shown as a lab demonstration by collecting a mason jar of lake or marsh mud with some water on top and exposed to the air. Add some ferric chloride or other iron salt. Typically a rusty looking flocculent layer will form within a few days on top of the sediments. Then, add a few millimeter thick layer of mineral oil or vegetable oil to seal the water off from atmospheric oxygen. After a few days to a week, depending on the organic content of the mud, the system will become anoxic, and the floc will dissolve. Note: Ferric chloride is also added to secondary treated wastewater (i.e., the organic matter has been largely decomposed but the nutrients remain) at some advanced wastewater treatment plants to precipitate phosphorus. The floc adsorbs phosphate, and then the water is filtered through sand or settled to remove the floc. It will also remove particles of silt and can be used to clarify water for both drinking water and wastewater treatment. anoxia

Sulfide and iron – summer vertical profiles Oligotrophic Eutrophic H2S O2 T O2 T Soluble Fe H2S Soluble Fe Depth anoxia Under oxic conditions, iron will be in the oxidized, ferric state, which immediately forms insoluble ferric hydroxide complexes. These compounds agglomerate to form relatively heavy, rust-like flocculent particles that rapidly settle out of the water column. Therefore, most of the year there is virtually no soluble iron in the upper water column. If bottom water becomes anoxic (due to bacterial respiration), ferric iron that is settling, as well as the material that is in the surficial layer of the sediments, can be chemically reduced to ferrous iron. This stuff is soluble in water, and, because it is at a much higher concentration at the sediment-water interface, it diffuses upwards. In doing so it also releases phosphate that was adsorbed to it. As the season progresses, organic matter continues to “rain” down, stimulating bacterial activity even more. Conditions in the deeper water become more and more chemically reducing as different groups of anaerobic bacteria consume available electron acceptors. After all the oxygen and nitrate are removed, a group of obligate (meaning oxygen is toxic to them at all times) anaerobic bacteria can use sulfate (SO4-2) as a terminal electron acceptor during their metabolism of organic matter (carbon). These are heterotrophic bacteria that use sulfate as they oxidize organic carbon to yield energy while releasing hydrogen sulfide (gas that smells like rotten eggs) as a by–product. Animals, such as ourselves, use a similar process but use oxygen as the terminal electron acceptor in respiration – oxidizing carbon to yield carbon dioxide and energy and reducing oxygen to water vapor, which we exhale along with the CO2. anoxia

18 4 4