Article Water Chemistry

In examining the physical properties of water, we have, in a way, begun discussing the chemical properties of water. We have already seen that H2O is a polar molecule with a slight positive charge near the hydrogen and a slight negative charge near the oxygen. These subtle charges are the key to understanding most of water's chemistry. We have seen that polar and
ionic molecules can dissolve in water, while ionically balanced compounds, including many hydrocarbons, are insoluble in water. If these chemicals are on the surface of a solid, they will determine whether that solid is hydrophobic or hydrophilic. Overall, however, the chemistry of water is largely a study of what is dissolved in it, and we will turn to those matters now.
Physics
Habitats
Excretion
pH
Pure distilled water is not a homogeneous mixture of H2O molecules. Water molecules tend to ionize, that is to separate into H+ and
OH- ions, which mix with the water. Normally, there are relatively few of these ions, and it is customary to measure the number of H+
ions present to get a value known as pH. In pure H2O at room temperature (20o C), there are about 0.0000001 grams of H+ ions per
liter (and an equivalent amount of OH- ions). You can see that 0.0000001 = 1 x 10-7, and, dredging up memories of mathematics long
forgotten, you may even recall that the logarithm of this number is the exponent (-7). Make that exponent positive (take the negative
exponent) and you have 7, the pH of pure water at room temperature.
Acids are substances which decrease the pH of a solution. They work by splitting up water molecules, moving in with the OH-, and
leaving "orphan" H+ floating around. Bases work the opposite way; they split up the H2O, but they combine with the H+ and leave
extra OH- in solution. Of course, there are also those acids and bases which affect pH directly; hydrochloric acid (HCl) in H2O
Most organisms operate best when the pH surrounding their cells is about 7, and natural waters usually oblige them. The pH of
clean mountain lakes (almost pure water) is usually about 7; the pH of the oceans is about 8. Local conditions can affect pH. Addition
of CO2 to water decreases pH (makes the water acidic); removal of CO2 makes water basic. Respiration and photosynthesis can
therefore have an effect on pH, particularly in areas where water flow is negligible.
Buffers are compounds which reduce the effects of acids or bases. The presence of a buffer usually prevents major changes in pH
even if a strong acid or base is added to the system. In the human body there are numerous buffer systems to prevent pH changes;
these include proteins and inorganic buffers. In natural systems, all of the buffers are inorganic. Sodium bicarbonate (NaHCO3) is one
of the major natural buffers; it is abundant in the ocean, and in freshwater where there are limestone rocks (which are largely
carbonate deposited from earlier, buffered oceans). Areas such as the Adirondacks, with mostly granite rocks and little limestone,
have little buffering capacity and thus are susceptible to pH changes due to acid rain. The "buffers" in things like aspirin (an acid) prevent stomach pH from becoming even more acidic.
The pH of a solution is critical since it changes a number of parameters. Some substances change their solubility or reactivity as pH changes. Proteins may change their shape and thus change their reactivity. Almost anything that dissolves in H2O will affect the pH somewhat, and we will consider those effects as we examine those substances.
Salinity
When water falls to earth, it is chemically pure - well, at least it used to be. In any event, it normally carries relatively few dissolved chemicals. We say that it has a low salinity. As it flows towards the oceans, it dissolves more and more molecules of various sorts from the rocks and soil it passes over - a process we call weathering. It also picks up materials released from plants and animals, either as waste products or through decay of their bodies. Since water does not leave the oceans except by evaporation (which
removes only water, leaving the other materials behind), these materials accumulate in the oceans, and the oceans have a salinity much greater than that of freshwater. In turn, the salinity of freshwater increases as it flows to the oceans, and coastal areas such as estuaries may have a salinity reflecting the mixture of fresh and salt water. Other areas that allow water to leave only by evaporation also may also have salt water. For instance, if you do not periodically allow water to flow completely through your houseplant pots, or if you never remove water from your freshwater aquarium, they will become salty. We call it salt water because most of the dissolved substances are just that - salts, simple ionic compounds that typically disassociate completely into positive and negative ions in water. Since an equal number of positive and negative ions are added in such cases, most salts do not change the pH appreciably, unless one of the components happens to be H+ or OH-. Note, however, that the presence of salts may affect the chemistry of certain pH tests, which may only be accurate when used in either fresh or salt water (some chemical tests will work in both). The most abundant of the salts found in the oceans is NaCl, sodium chloride or table salt. We measure salinity in terms of the number of grams of dissolved salts in 1000 g (one l) of seawater. Since all these salts affect the electrical conductivity of the water, it is simple to measure salinity with an electrical meter. The more current that passes through a solution, the greater its salinity. Seawater ranges in salinity, but a useful approximation is 35 g / kg; or 35 parts per thousand or 3.5%. Places like the Great Salt Lake, certain tidal pools, etc., can have higher salinities; most freshwater systems have dramatically lower
salinities. In seawater, the ions are dominated by Cl- (19.353 gr/kg), Na+ (10.76 gr/kg), SO4-2 (2.712 gr/kg), Mg+2 (1.294 gr/kg), Ca+2 (0.413
gr/kg), K+ (0.387 gr/kg), HCO3- (0.142 gr/kg), Br- (0.067 gr/kg), and Sr+2 (0.008 gr/kg); other ions are present in trace amounts, including gold. Freshwater is dominated by similar ions, but the amounts are highly variable and depend on season, amount of rainfall, type of rocks, etc. Freshwater with an abundance of calcium and magnesium is termed hard; it resists forming suds when detergents are added and leaves a residue. Water softening agents often attempt to chelate these ions out of solution.
The salinity of water is important to organisms for two main reasons: osmoticity and density. We have already considered density;
remember that salt water is more dense than freshwater, thus, it is easier to float in salt water. This also means that freshwater flowing into the ocean tends to remain on the surface for some time, and, further, that it takes some energy (tides, wind, currents) to get the two to mix effectively. It is not uncommon in estuaries for freshwater to lie on top of saltwater in layers. Temperature still plays a role; warm freshwater is very likely to lay on top of cold seawater, but other possibilities may also occur, such as warm seawater
laying on top of a layer of cold freshwater. We will take up the concept of osmoticity next.
Osmotic Relations
Imagine water of two different salinities. We say that the more saline solution is hypertonic (or hyperosmotic) in relation to the other,
or that the less saline solution is hypotonic (or hypoosmotic) in relation to the other. If the salinities were the same, they would be isotonic (isosmotic). Osmoticity, then, is simply a comparison of the salinities (or, more accurately, the number of dissolved particles, including non-ionic compounds) of two solutions. It is important to remember that, in considering osmoticity, that you must have two different solutions to compare - it is a relative term.
If you mix the two solutions, they will achieve an intermediate salinity. The process of diffusion means that all of the ions will, as a result of random chance, distribute themselves evenly throughout the solution. This occurs because the molecules and ions are in constant motion (except at 0 K); diffusion occurs more rapidly at higher temperatures. Stirring, mixing, bulk flow due to density differences (as a result of temperature or salinity) etc., will all decrease the time it takes for the solutions to come into equilibrium, with the dissolved material equally distributed.
Now, place the two solutions in contact with each other, but separated by a membrane. The membrane will have the property of allowing only small, uncharged (polar is OK too) molecules to pass through. Say the one solution is in a sack made of the membrane, and the sack is dropped into the other solution. Will diffusion take place? The answer is yes and no. Water molecules, being small and (relatively) uncharged, pass through easily, and in both directions. Ions get stopped at the border due to their charge, as do larger molecules such as sugars and proteins, regardless of their charges. For
physical reasons we need not get into here, the water molecules will, on average, tend to go more to the hyperosmotic side of the membrane, although a few rugged individualists will still cross over from the hypertonic side to the hypotonic side. The net flow of water will be to the hypertonic side. If the outside solution in our example is hypertonic, water will flow out of the sack and the sack will shrink; if the outside solution is hypotonic, water will flow into the bag and it will swell and (perhaps) burst. We have just built a model of a cell (Fig. 4)(ions can penetrate the membrane of a cell, however)known as osmoconformation, and works only if the organism is in an unchanging environment - like the middle of the ocean, or deep inside a body. It is no accident that the salinity (though not the precise ionic makeup) of our cells - and the cells of most living things - is isotonic with seawater.
Moving into freshwater (or moving onto land) requires a different strategy. Freshwater is simply too dilute to keep life in a cell going. You need more stuff in your cells, and that stuff increases tonicity. Freshwater organisms face a constant influx of water from the surrounding hypotonic medium, and they can potentially lose important ions to that solution also. Therefore, the strategy among most freshwater organisms is to cover as much of the body as possible with an impermeable coat, and leave all water exchange to a relatively small number of cells. These cells will maintain the water balance, and the remaining cells are bathed in an isotonic solution. Cells can maintain osmotic balance by using ATP to actively pump Cl- ions into the cell. The inside of the cell becomes negatively charged, and other ions, such as Na+ come in because of this. Water that flows in is passed on to the blood and excreted as a dilute urine.
Osmotic exchange also takes place across the lining of the gut. The freshwater fish gains water but loses ions passively across the gills; to compensate, the gills actively pump in ions and the kidneys form a dilute urine. The saltwater fish gains ions and loses water across the gills; to compensate, water is ingested (along with salt), the gills actively pump ions out of the body, and a small amount of relatively concentrated urine is formed. Many marine organisms (those which maintain their internal fluids hypotonic to seawater, perhaps reflecting the salinity of the sea when they evolved), and all terrestrial organisms, face the problem of water loss. This is conquered in reverse fashion. These organisms drink seawater, absorb water (and ions) from the gut, and pump ions out through specialized cells. The specialized cells in marine organisms are the same as those used by freshwater organisms to pump ions in; it's just that the membrane-bound proteins which form the ion "pumps" are "installed" backwards. These Cl- pumping cells, whichever way they pump, are called chloride cells. In fish (both freshwater and marine), they are located on the gills. Because respiratory structures must have relatively permeable surfaces for gas exchange, they are also a common place to put chloride cells on a body which is otherwise impervious to water flow. Another popular place is in the gut and kidneys, in both places ion concentrations are manipulated to get water to flow where the organism wants it to.
Dissolved Gasses
The atmosphere is about 78% nitrogen, 21% oxygen, and only 0.033% carbon dioxide (but we're working on that). These three gasses have different solubilities in H2O, however. To measure solubility, imagine this experiment: Replace all the air in the atmosphere with only the gas you are interested in so that at ground level the pressure is 1 atmosphere. If you now take one liter of water (let's say at 10o C), it will become saturated with the gas. In our imaginary world, if the gas was nitrogen, the water would hold 18.61 ml of it; if the gas was oxygen, the water would hold 38.46 ml of it; if the gas was carbon dioxide, the water would hold a whopping 1,194 ml of it! If you're wondering how 1 liter of water can hold 1.194 liters of gas, get a liter of Perrier and shake it up - then open it. The solubility experiment tells us that oxygen is about twice as soluble as nitrogen, and that CO2 is about 1,000 times more soluble than either nitrogen or oxygen. Of course, in the real world there is not 1 atmosphere of pressure of each of these gasses, each gas has a partial pressure corresponding to the percentage of the atmosphere that it makes up. Multiplying the solubility of the gas by its percentage in the atmosphere gives the amount of that gas that will be dissolved in the water:
Solubility of Gasses in H2O at 10o C
Gas Percent Solubility* In water*
Nitrogen 78.084% 18.61 14.53
Oxygen 20.946% 38.46 8.06
Carbon Dioxide 0.033% 1,194.00 0.39
* Solubility in ml/l
The absolute amount of a gas in water solution will vary with several factors. Increasing temperature will reduce the amount of gas that water can hold; you are familiar with this fact already, since it is manifested whenever you heat water (the small bubbles that form before the water boils). Decreasing pressure (increased altitude) will also decrease the amount of gas dissolved. Increasing salinity also decreases the ability of water to dissolve gasses; seawater holds about 20% less gas than freshwater, and hypersaline water
holds even less gas. And, of course, there are other gasses which are dissolved in water besides these three (which are the major ones).
Oxygen:
Oxygen, of course, is of critical importance to living organisms. It is important to remember that oxygen is a potent poison, and that too much can be a bad thing. We're probably lucky that the atmosphere has only 21% O2; astronomers on other planets have probably concluded that life cannot exist on Earth because of the high levels of noxious oxygen. Any more O2 and it would probably start to poison the process of photosynthesis (actually, it already does to a point, the reaction is called photorespiration). In water,
however, it is easy for O2 concentrations to become depleted in local areas, such as in sediments or the bottom of stagnant bodies of water. These local areas of oxygen depletion serve as important refugia for the many ecologically significant anaerobic organisms, and thus should not be viewed automatically as "bad". Nitrogen fixation, for instance, is notoriously sensitive to O2; root nodules of legumes and heterocysts of Cyanobacteria are structures built to exclude O2 so that nitrogen fixation can take place.
The solubility of O2 in water is slightly more complex than was explained initially above. Better estimates of O2 solubility involve moist air as opposed to the dry air in our theoretical planet experiment. Oxygen solubility is also very temperature dependent. In general, the solubility of oxygen can be estimated by the equation:

mg/l=468/31.6+t
Where t is degrees C; note that we are now talking about milligrams, not milliliters. Thus, for a t of 10o C, there should be 11.25 mg O2/l; at 20o C there should be 9.069 mg O2/l; and at 0o C there should be 14.81 mg O2/l at saturation. These values are for sea level and fresh water; at altitude or in salt water the saturation values would be lower. Water is not always at saturation, however. Organisms within the water can quickly use up dissolved O2, and replacement by diffusion may not equal the rate of respiration, thus
resulting in lower O2 levels. Oxygen enters the water through diffusion from the atmosphere and from photosynthesis. Diffusion from the atmosphere may be enhanced by any type of turbulence such as water flowing over rocks, waves, wind, etc.; photosynthesis is dependent on many factors. Oxygen loss is primarily due to respiration of animals and plants in the water; respiration is temperature dependent, with more respiration occurring at higher temperatures. These relationships set the stage for determining the critical factors for O2 availability in water. More oxygen will be available in cool,
sunny, turbulent habitats (uptake of O2 by organisms is also enhanced by flow or turbulence since more water will pass over
respiratory structures per unit time). Oxygen will be in short supply under warm, dark, stagnant conditions, where respiration is high and no photosynthesis is taking place. In natural aquatic systems, O2 stress most often occurs at the water-sediment interface on dark, calm nights in the middle of the summer or dry season, when water flow is at a minimum. Any pollutant which adds to the amount of plant life present (fertilizers), or which contributes organic matter which will be decomposed by aerobic bacteria (sewage),
will aggravate the O2 stress at that time. Also in natural systems, the most O2 is available in cool habitats with shallow water, particularly in turbulent headwater streams or near wave-swept coasts. Oxygen concentration in the water can be measured by either electrical meters or by chemical methods. The meters use electrodes and must be calibrated or adjusted to compensate for temperature (most include an electrical thermometer attached to the probe), altitude, and salinity. Chemical tests, such as the well-known Winkler Method, which measures oxygen by determining how much of a standard chemical is oxidized, work by titration. Results are often expressed in percent saturation; for instance, if the water is at 20o C and measures 8 mg O2/l rather than the saturation value of 9.069 mg O2/l, the percent saturation would be 88%. It is not unknown to have supersaturated water; this results in percent saturations over 100%, under such conditions O2 is released to the atmosphere - which is how it got there in the first place!
Another important concept to consider is Biochemical Oxygen Demand, or BOD. BOD is the measure of how much oxygen is taken up from the water by both biological agents (organisms) and simple chemical reactions (like Fe + O2 = rust). To measure BOD, a sample of the water (and/or sediment) is measured for O2 content, sealed up, and left for a specified period (often 24 hours). The O2 level is measured again, and the amount of O2 used up (mg), or the rate of O2 use (mg/hour) is reported as the BOD. Regulations governing sewage plants, stockyards, etc., may specify the amount of BOD that can be released each day.
Carbon Dioxide:
Carbon dioxide is ying to oxygen's yang. As one is evolved in respiration or photosynthesis, the other is taken up. Still, the much greater solubility of CO2 compared to O2 means that the relationship of CO2 to water is much different compared to that of O2 and water. For one, CO2 dissolved in H2O increases acidity. Secondly, CO2 can combine with other chemicals in the water to form different compounds; these compounds further affect the chemistry of the water. The primary sources of CO2 in water are biological
(respiration), geologic (weathering of carbonate containing rock such as limestone) and from the atmosphere. Rainwater in particular, with its high surface area, picks up a lot of CO2 as it falls. A small amount of CO2 entering water combines with the water to form carbonic acid:
CO2 + H2O <--> H2CO3
The carbonic acid, in turn, may dissociate:
H2CO3 <-->HCO3- + H+
The HCO3- is known as bicarbonate, and may be taken up by photosynthesizing plants. Changing the pH of a solution changes the
relative concentration of the various CO2 forms; a decrease in pH (more acid) leads to more CO2 and H2CO3-; an increase in pH (more basic) leads to an increase of the CO3-2 ion. In order to measure the amount of CO2 present, one must obviously know the pH of the solution. Carbon dioxide may also combine with water and metals such as magnesium and calcium to form other bicarbonates. The amount of CO2 so combined is referred to as alkalinity, which really has nothing to do with OH- concentration, but much to do with the buffering capacity of the water. It works like this: Highly alkaline water tends to have a high (basic) pH and will turn a phenolphthalein solution
pink. If you add acid to it, the bicarbonates, with their negative charge, attract and bind the positive H+ ions, and form carbonic acid. If you keep adding acid, eventually the pH changes to 8.3, and the pink fades. The amount of acid added corresponds to the phenolphthalein alkalinity, but not all the bicarbonate is converted at this point; in fact, it is at its peak. If you now add methyl orange, a dye that will change color at pH 4.4, and continue to add acid, you will drive more bicarbonate to form carbonic acid, which in turn reaches its peak at a pH of 4.4. The total amount of acid added thus corresponds to the amount of CO2 present in the sample. This method works only if there are not significant numbers of non-carbonate negative ions to absorb H+ ions. Testing for CO2 is usually done chemically, and is quite complicated. It is based on the above reactions and works something like
this: if phenolphthalein added to a water solution turns pink, then the pH is over 8.3 and significant bicarbonates, including those of Ca++ and Mg++, are present. The solution is titrated with acid to a pH of 4.4, and the total amount of acid added corresponds to the amount of CO2 present (as bicarbonate). If the initial solution with phenolphthalein does not turn pink, then a lesser amount of CO2 (as bicarbonate or whatever) is present, and bicarbonate is titrated until a pink color appears. The amount of bicarbonate added will then correspond inversely to the amount of CO2 that was present initially; the more bicarbonate you have to add, the less was there to start with. An interesting reaction may take place when CO2 is removed from water during photosynthesis - a precipitate of CaCO3 (calcium carbonate) may form. This reaction, along with calcium carbonate formation by corals, has lead to the production of most of the limestone in the world. Weathering of that limestone, in turn, yields much of the carbonate present in freshwaters, and, as we have seen, that carbonate forms an effective buffer against decreased pH due to acid rain. As a result of the many forms it can take, CO2 is usually present in ample amounts for photosynthesis to occur in aquatic habitats. Although rarely a limiting factor, it is known that more alkaline lakes are, up to a point, more productive than less alkaline lakes. However, as we shall soon see, other factors, such as the abundance of phosphorous and nitrogen, are more often limiting factors, and levels of these plant nutrients are usually correlated with alkalinity, making the independent factors hard to sort out.
Plant Nutrients - Nitrogen and Phosphorous
Aside from O2 and CO2, there are a variety of other chemicals needed by living things. Animals, almost by definition, obtain these other chemicals along with the carbohydrates and proteins they ingest when they consume other animals or plants, therefore, animals are usually content as long as the water has enough O2 and a decent salinity. Plants, on the other hand, are more self-sufficient, and they can synthesize a wide variety of complex molecules from simple inorganic precursors. The ecological community that will develop in a body of water is thus often dependent on the suitability of the habitat for the growth of photosynthetic organisms. Exceptions to this include the deep ocean (which is dependent, however, on the growth of phytoplankton above), or headwater streams (which depend on adjacent trees for most of their organic input through leaffall), or cave streams (bat guano) and so on.Along with sunlight and CO2, the major needs of a plant include macronutrients nitrogen and phosphorous (used for proteins, DNA,
RNA, ATP, etc.) and micronutrients such as sulfur (protein), magnesium (chlorophyll), and iron (cytochromes) (this list is not comprehensive). The micronutrients may be found in very small concentrations; plants (I will use the term plants to refer to all photosynthetic organisms) are good at obtaining them even if they are in low concentration; and they are rarely a limiting factor (be sure to review your ecology to be sure that you understand the concept of limiting factor). Some micronutrients, such as sulfur (as SO4-2), magnesium (as Mg+2), calcium ( as Ca+2) and potassium (as K+) are important constituents of both seawater and freshwater, as we have seen earlier. Phosphorous (P) and nitrogen (N) are critical to plant growth, and they (usually P, but sometimes N) are often limiting factors to plant growth. Before you object, recalling that dissolved nitrogen is common in water, remember that it is dissolved nitrogen gas, N2, which is inert and cannot be used by most plants. The exception here are the cyanobacteria, which can fix N2 in the heterocysts, which provide a local anoxic environment for the nitrogen-fixing enzymes (and bacteria in anoxic root nodules of legumes and other anoxic places in the soil). Bodies of water with a low N/P ratio are thus prone to blooms of cyanobacteria. For most plants, N must be in the form of nitrate (NO3-) or ammonia (NH3, NH4+ in water). Ammonia, of course, is the nitrogenous waste of choice for many aquatic organisms, and even more is released by bacteria breaking down dead plants and animals or other nitrogenous animal wastes such as urea. Nitrate is a product of the nitrogen cycle; the nitrogen cycle in water differs slightly from the nitrogen cycle that takes place on land (which you are probably familiar with). On land, N2 is fixed by bacteria in the soil such as Rhizobium, Clostridium, and Azotobacter; in the water (both freshwater and marine) N2 is fixed by such cyanobacteria as Anabaena, Plectonema, and Nostoc. The reaction requires energy and proceeds as follows:
Nitrogen Fixation: 2N2 + 6H2O ---> 4 NH3 + 3O2
Ammonia, whether generated by nitrogen fixation or by the breakdown of amino acids by animals or decomposers, is toxic. As the pH of water increases, more of the ammonia exists in the water as NH4+. NH4+ is even more toxic than NH3, and the fact that it is more prevalent at higher pH leads to one of the significant differences between keeping a marine and a freshwater aquarium. A marine aquarium typically has a pH of 8.0 to 8.5; a freshwater aquarium will usually have a pH of about 7. At pH 8.0, there is far more NH4+ present, and, if too many animals are producing too much NH3, then NH4+ levels will soon become toxic. Therefore, marine tanks must be "aged", that is, stocked slowly, to allow populations of bacteria to develop to remove the ammonia. The role of these latter bacteria will be explored below.
Once produced, ammonia (NH3) is used by a variety of plants and bacteria as the source of the amino group for amino acid synthesis (another reason that freshwater tanks are more tolerant than marine tanks in regards to ammonia is the ready availability of
freshwater plants which help reduce ammonia levels). The amino acid synthesis reaction also requires energy and looks like this:
Amino acid synthesis: 2NH3 + 2H2O + 4CO2 ---> 2 CH2NH2COOH + 3O2
Note that both this reaction and the preceding one release O2 into the atmosphere; photosynthesis is not the only source of O2 in the atmosphere! While amino acid synthesis does remove some ammonia from the water, much more is usually present. Another reaction, nitrification, takes ammonia and converts it to nitrite (NO2-); this reaction releases energy to the organism which carries it out:
Nitrification I: 2NH4+ + 3O2 ---> 2NO2- + 4H+ + 2H2O
In water, this reaction is carried out mostly by bacteria of the genus Nitrosomonas. Nitrite is less toxic than ammonia, but is still toxic; high levels of nitrite can kill many aquatic organisms. Fortunately, a further nitrification reaction can occur (also with a release of
energy):
Nitrification II: 2NO2- + O2 ---> 2NO3-
The end product here, nitrate (NO3-), is even less toxic than nitrite, and can be used by many plants as a nitrogen source. In aquatic systems and terrestrial systems as well, this reaction is carried out by bacteria of the genus Nitrobacter. In a typical marine aquarium, nitrate may approach toxic levels, but this process takes months. In addition, a number of denitrification reactions take place and reduce nitrate levels, as does uptake by plants.
To retrace the nitrogen cycle, let us consider the marine aquarium again (Fig. 7). Ammonia levels build as animals excrete nitrogenous wastes; as they die and decompose; as food (with protein) is added; and as N2 from the atmosphere is fixed by cyanobacteria. Because of the high pH (8.0), most of the ammonia will exist as toxic NH4+. Nitrosomonas bacteria will convert the
ammonia to nitrite, and Nitrobacter bacteria will convert the nitrite to nitrate, which can be utilized by plants. Denitrification will remove some of the nitrate from the water. In a freshly established marine tank with a few fish, it is not uncommon for the ammonia levels to
peak, then drop as the Nitrosomonas bacteria take hold and begin to convert ammonia to nitrite. As nitrite levels build and peak, Nitrobacter populations will thrive and convert the nitrite to nitrate, reducing nitrite concentrations to near zero. Typically, it takes about one month for the bacterial populations to become established, and it is usually wise to monitor the process by daily tests of ammonia and nitrite levels. The number of organisms that can be maintained in a marine tank is usually proportional to the amount of surface area on the gravel of the aquarium available for the Nitrosomonas and Nitrobacter bacteria to attach to.
N2 in the air is fixed by cyanobacteria and put in the form of protein, which is eaten by fish. Fish release NH3, which is taken up by heterotrophic bacteria, plants, cyanobacteria, and the bacterium Nitrosomonas. The first three reincorporate the NH3 into protein, Nitrosomonas converts it to NO2, which is taken up in turn by Nitrobacter, which gives off NO3. The NO3 is a plant nutrient and is also utilized by anaerobic bacteria, which can produce N2, NO2, or NH3.
The story with phosphorous is much simpler. Phosphorus is typically available to plants as organophosphate (PO4-3); this compound is a common weathering product of igneous rock. Other sources of phosphate are decaying animal bodies, animal wastes, bones of vertebrates, and bird guano. The last is mentioned specifically because it is such a rich source of phosphorous; in fact, huge quantities of bird guano that accumulate near nesting areas of shorebirds are often mined for fertilizer. Phosphorous trapped in the bodies of dead organisms which sink to the bottom may accumulate in bottom sediments; this phosphorous will become available again to living plants (animals do not take up phosphorous directly from the environment) when currents sweep the bottom and bring it to the surface where there is sufficient light for photosynthesis (remember the LCP?). In nature, two such currents are noteworthy, the overturn of lakes (which occurs when the water is all of the same temperature and density, more on this later), and upwellings, places where cold ocean currents meet continents and rise up from the bottom. The most famous example of an upwelling is off the coast of Peru; as the cold water reaches the surface (replacing warm surface waters blown west by prevailing winds), it brings up
phosphorous, which encourages the growth of abundant algae, which in turn are appetizers for anchovies, which are then fished in great numbers by humans (for pizza topping) and birds (where do you think the guano comes from?). Phosphorous is often the limiting factor for plant growth in aquatic systems. In many aquatic habitats, it is impossible to measure any free phosphate in the water when heavy algal growth (blooms) are in progress, mostly because the algae use up the phosphate as soon as it becomes available. Many algae store inclusions of phosphorous (when available) in their cells as a "hedge" against later phosphate shortages. Tremendous algal blooms are a symptom of eutrophication, a natural process that occurs as lakes age and accumulate phosphorous. Such huge blooms of algae can cause problems when they die and decompose, or when, particularly on hot, still summer nights, their respiration uses up all the O2 in the water. Many human activities accelerate the process of eutrophication; this leads to a phenomenon known as cultural eutrophication to set it apart from the natural process. In cultural eutrophication, phosphorous from agricultural runoff (phosphate fertilizers, animal wastes), human sewage, detergents, etc., is added to the water. The much publicized "death" of the Great Lakes was (is) due to cultural eutrophication. A eutrophied lake is typified by heavy algal growth, turbid water, and overall high productivity. The opposite, an
oligotrophic lake, has clear water with little life (paradoxically, highly eutrophic lakes such as our "dead" Great Lakes, are teeming with life, while pristine oligotrophic lakes are nearly sterile). Remedies to cultural eutrophication include soil erosion control (much phosphorus is carried into aquatic systems bound to soil particles), careful use of fertilizers (the use of fertilizers on grass is criminal and should be totally banned), and tertiary sewage treatment (primary treatment kills germs, secondary removes solids, tertiary removes phosphorus and other chemicals). Even after the source of phosphorus is found and diverted, the sediments may contain enough phosphorus to keep the lake eutrophic for years. Again, under natural conditions, oligotrophic lakes occur near headwaters where little weathering has occurred to release phosphorous, and eutrophic lakes are found downstream (downstream is usually geologically older), where phosphorous inputs from the upstream sites make abundant algal growth possible.


Rick