INTRODUCTION

BACKGROUND

The Early Ocean

The Modern Ocean

The Reef Aquarium

The Boggs Aquarium

EXPERIMENTS

Salinity

Alkalinity

Dissolved Oxygen

Sulfur -- Sulfate

Calcium

Magnesium

Phosphorous -- Phosphate

Nitrogen – Ammonia/Nitrite/Nitrate

CONCLUSION

The ocean environment is the origin of life on earth. Organic molecules formed in an organic "goo" composed of just simple elements: carbon, hydrogen, oxygen, nitrogen, phosphorous, and sulfur. This lab attempts to understand a part of the roles of these elements and several simple compounds or molecules in today's ocean environment. The subject of study in an artificial reef system set up in an 300 gallon aquarium in Boggs Chemistry and Biochemistry building.

Water is the only pure liquid on earth. It is a polar solvent and is responsible for life. It was in the water that the six main elements of life began to interact with the catalyst of potential energy from lightning. The atmosphere was composed up of mostly carbon dioxide (CO₂) and nitrogen (N₂) gases. One of the first large steps towards higher life forms was the production of oxygen (O₂), about 200 million years ago. This gas was produced through a process known as photosynthesis. Photosynthetic bacteria began to use the sun as their energy source. This solar energy was coupled with carbon and nitrogen from the atmosphere, to begin to form simple biochemical compounds. The O₂ product was not a atmospheric gas at first. It took over 100 million years before the oxygen produced by the ocean life was enough to react with the iron and other elements in the sea.

2Fe²⁺ + O₂ à 2FeO

2H₂S + O₂ à 2SO + 4H⁺(2H₂)

Once an equilibrium was established between the oxygen and these other chemicals in the water of the ocean, oxygen was finally released and began to build up in the water. When the water was at a maximum dissolved oxygen concentration, oxygen was also released into the atmosphere.

The next large step after photosynthesis was the beginning of nitrogen fixation. Nitrogen fixing bacteria were present in the water. Ammonia was made from the hydrogen produced from the oxygen's reaction with hydrogen sulfide and from the N₂ that was present in both the ocean an the atmosphere.

N₂ + 6H⁺ à 2NH₃

Ammonia collected in the sediment at the ocean floor. As the nitrogen cycle continued to develop, more and more chemical species were continuing to change the composition of the water.

The modern ocean is like one very complex and ever-changing living organism. There are many different aspects, including chemicals, chemical cycles, organisms, etc. that are required for its health.

Salinity is important to a marine ecosystem, because it is necessary to maintain constant conditions. Salinity includes several ions in solution including calcium, magnesium, potassium, and chloride. Natural saltwater environments have a salinity concentration of 35 grams per 1000 grams of water, or 35 ppt.

Salinity actually limits the amount of oxygen that can be dissolved in the water. As the salinity of the water increases, the amount of oxygen that is dissolved in the system decreases. Changes in the salinity of the system can be a result of evaporation. Evaporation increases the concentration of the chemical ions, so as water is evaporated, the salinity increases and the concentration of dissolved oxygen decreases. Salinity must be monitored in order to maintain a constant and healthy level of dissolved oxygen.

Alkalinity in a system is a measurement of hydroxide ions, carbonate ions, and carbonic acid. It is sometimes referred to carbonate hardness, or KH, which is specifically the concentration of calcium carbonate (CaCO₃). General hardness is, however, the amount of calcium and magnesium. The alkaline nature of the aquarium acts as a natural buffering system. An aquarium system maintains a fairly constant pH of approximately 8.3, while a system low in alkalinity has an unstable more acidic pH.

Carbon dioxide also plays a role in the alkalinity of a system. Carbon dioxide may come from several sources including the following: exchange with the atmosphere and the surface, runoff, bacteria, and aquatic animals during respiration.

CO₂ + H₂O l H₂CO₃

Oxygen is an essential component of all living things. Within the marine ecosystem oxygen is recycled through respiration and photosynthesis processes just as it is the terrestrial environments. Animals take dissolved oxygen directly from the water and produce carbon dioxide. Carbon dioxide is consumed by plants during photosynthesis to produce food for both plants and animals. There are several sources of oxygen into the system. Oxygen may enter the system through exchange at the surface of the water. It make sense then that water with more surface movement has a higher dissolved oxygen concentration. The normal level of dissolved oxygen in a marine environment is between 4 and 6 parts per million, while in a freshwater system the level is closer to 8 parts per million.

Sulfate (SO₄^2-)is a very common polyatomic ion that is present in the aquatic marine environment. Sulfate is actually one of the most common metal salts, along with nitrate. Most sulfates are highly soluble except for lead(II) sulfate (PbSO₄) and BaSO₄. (This property was the primary characteristic of sulfate used to determine the concentration of sulfate in the water and will be discussed more later.) The SO₄^2- ion is neither oxidizing nor reducing. The SO₄^2- ion is involved in an acid/base system with the hydrogen sulfate (H₂SO₄) and the hydrogen sulfide ion (HSO₄^-), although the presence of SO₄^2- does not noticeably alter the pH of the system. Also, sulfate creates very thermally stable compounds.

The exact role of sulfur in the water chemistry of the aquatic environment is not fully understood. Sulfur does have a role in the organic part of the ecosystem. It is a major building unit in several amino acids including: cysteine, cystine, methionine, homocysteine, cystathione, glutathione, ctsteic acid, taurine, biotin, and thiamine. There is also a sulfur cycle within the environment. Chemosynthetic bacteria oxidize sulfur forming sulfate:

2S + 2H₂O + 3O₂ à 2H₂SO₄(aq).

The chemosynthetic bacteria oxidize the sulfide to get energy needed to extract the carbon from carbon dioxide (CO₂) in the water. This carbon is used to build the organic molecules required for life.

The sulfur cycle, starting with native sulfur, is oxidized by Thiobacillus which produce sulfate. This sulfate reacts through synthesis by microbes and plants producing organic sulfur, but sulfate is also produced from organic sulfur by oxidation by microbes and animals. The organic sulfur is transformed to sulfide by the action of heterotrophic microbes. Sulfide is then oxidized to produce native sulfur by oxidation by several types of bacteria, especially Thiothrix and Beggiatoa. There is also an exchange between sulfate ad sulfide through oxidation and reduction reactions.

Specific areas of the ocean are found to be very high in sulfur levels. The Galapagos Rift, which is about 2.5 kilometers deep, 100 degrees Celsius, and at a pressure of 280 kilograms per square centimeter, is one location of chemosynthetic bacteria which reduce sulfate to produce H₂S as their source of energy. The Palos Verde peninsula is another location of toxic sulfur levels. Sulfur escapes through fissures in the sea floor, very shallow at just 1 to 10 meters deep. Chemosynthetic bacteria surround these fissures to obtain some of the sulfur nutrient.

Calcium (Ca²⁺) and magnesium (Mg²⁺) are the fifth and eighth most abundant elements in the earth's crust respectively. Study of Ca²⁺ and Mg²⁺ in the marine environment has sometimes been overlooked due to their weak interactions and reversibility of reactions. Calcium and magnesium are found in seawater as either free ions, Ca²⁺ and Mg²⁺, or loosely bonded to small organic molecules and are sometimes referred to as the general hardness of the water.

Some alkali earth metal cations have been found to bond with dissolved ATP molecules. Magnesium also plays a large role in the process of photosynthesis in plants. The central element in the chlorophyll molecule is magnesium. Mg²⁺ is the third most common ion in seawater, and this ion in seawater is the major source of magnesium metal for industry. Calcium is a major constituent in shells and skeletons.

It is also one of the necessary components of seawater. It is directly linked to plant growth, and it is essential in many biological processes where the native source of the phosphate is the seawater and sediment. Orthophosphate is a term used to describe phosphates in solution. Orthophosphates included PO₄^3-, HPO₄^2-, H₂PO₄^-, and H₃PO₄. The type of phosphate is determined according to the pH of the solution, more H₃PO₄ being present in a more acidic solution.

Phosphates have a specific cycle in the environment. Beginning in the soil, the phosphates are transferred into the water through drainage and runoff. These phosphates are either taken into plant and algal tissues or lost into deeper sediments Phosphates move through the cycle in shells and skeletons, animal tissues and feces, urine, and via decomposers such as bacteria and fungi. These phosphates are eventually returned back into solution. The phosphates in the soil may also be taken up by terrestrial plants. Both land animals and decomposers obtain nutrients, including phosphates, from these plants. The phosphates continue through the cycle by returning to the soil by way of urine, feces, tissues, and the result of decomposition.

Phosphorous is necessary for life: it is a macronutrient. Phosphorous is present in many organic molecules. P₂O₅ makes up 28% of deoxyribonucleic acid (DNA). It is also critical for metabolism and photosynthesis. Energy carrying molecules, adenosine triphosphate (ATP), require phosphate as part of their basic structure. Industry uses phosphorous to make fertilizers, to clean oil spills, and to improve food by providing calcium, potassium, and phosphorous into the diet. In meats, poultry, and seafood, phosphorus retains proteins and improves the color, texture, and flavor of the food.

In the ocean nitrate concentrations are greater deeper in the ocean. Nitrogen as well as phosphorous and silicon are quickly removed from the upper layers of the hydrosphere during photosynthesis. Phytoplankton and microscopic algae use the nitrate. Nitrogen and phosphorous exist in a relatively constant ration in the ocean, 16:1. In deeper waters, reoxidaton, or respiration, of nitrogen and phosphorous produce inorganic nutrients in the water. Also, conditions in the North Pacific would suggest that as levels of dissolved oxygen are decreased, the concentration of nitrate increases.

Actual levels for the marine environment should be very low, below 1 ppm. Nitrogen is present in the soil and in organic matter, but elevated levels of nitrogen, in ammonia, ammonium, nitrite, or nitrate form, signal a problem with the tank. Excess ammonia is toxic and has a negative impact on the ecosystem. In a tank the nitrogen levels must be monitored very carefully so that a constant and viable equilibrium can be maintained.

There are several ways to report nitrogen levels. Many results are listed as parts per million of the specific ion being tested, such as ammonia (NH₃), ammonium (NH₄⁺), nitrite (NO₂^-), or nitrate (NO₃^-). The true value of nitrogen in each test is actually determined by taking into account the other elements in the species being tested. These results are reported as N~NH₃, N~NH₄⁺, N~NO₂^-, or N~NO₃^-. The conversion between the concentration of the ions and the concentration of nitrogen are as follows:

[N~NH₃] = 0.8 [NH₃]

[N~NO₂^-] = 0.3 [NO₂^-]

[N~NO₃] = 0.23 [NO₃^-].

Ammonia and ammonium exist in the solution in equilibrium with the ratio depending on the pH of the system. The concentration of NH₃ increases with an increase of pH and an increase in temperature. The ammonia concentration is also more dependent on pH rather than temperature. An increase in the pH by one unit, increases the ammonia concentration by ten times. Normal ammonia levels are around 0.1 ppm. The concentration of NO₃^- is harmful to fish above a 60 ppm concentration. Nitrate is generally harmless to the system although is increases over time.

Nitrogen is the most abundant element present in the earth's atmosphere. It makes up seventy-eight percent of all atmospheric gases. The nitrogen in the atmosphere is in the form of dinitrogen (N₂). Living organisms are not able to utilize this elemental form of nitrogen. In order for living organisms to obtain this essential element, the nitrogen must be fixed in small molecules or ions such as nitrogen oxide (NO), ammonia (NH₃), ammonium (NH₄⁺), nitrite (NO₂^-), and nitrate (NO₃^-). This nitrogen fixation may be achieved with the aid of bacteria in a symbiotic relationship with plants or by lightning which converts dinitrogen and dioxygen into nitrous oxide:

N₂(g) + O₂(g) à 2NO(g).

Nitrogen is cycled through the environment, through several fixation states. This nitrogen cycle is composed of three main stages: ammonification, nitrification, and assimilation. Nitrogen is introduced into the environment via volcanic action. Volcanoes would spew ammonia (NH₃) or ammonium (NH₄⁺) into the atmosphere. These species enter the stage of nitrification. Bacteria, the Nitrosomonas bacterial species, in the soil oxidize NH₃ or NH₄⁺ to produce nitrite (NO₂^-):

2NH₃ + 3O₂(g) à 2NO₂^- + 2H⁺ + 2H₂O

2NH₄⁺ + 3O₂(g) à 2NO₂^- + 4H⁺ + 2H₂O.

This reaction is energy producing, and the bacteria that oxidize the NH₃ or NH₄⁺ use this energy as their primary energy source. Another type of bacteria, the Nitrobacter species, obtains energy through the oxidation of the nitrite ion to a nitrate ion (NO₃^-):

2NO₂^- + O₂(g) à 2NO₃^-.

Nitrate is the most common form of nitrogen in the soil, and it may take several different routes through the nitrogen cycle.

One route for nitrate is absorption into plants. The acquisition of nitrate through the root system is the main way that nitrogen is taken into the plant's system. Plants may also use ammonium directly, but nitrate is generally preferred. Any nitrate taken into the plant system is reduced back to ammonium. This process is called assimilation and requires energy. The plants use energy obtain through photosynthesis. The ammonia is then included in organic compounds such as the vital proteins, amino acids, nucleic acids, and nucleotides. As the plants die, the organic matter is reintroduced into the soil. Bacteria and fungi then decompose the material again to ammonia and ammonium to restart the ammonification process. Plants may also be eaten by animal life. Animals may either recycle the nitrogen through the excretion of urine which contains ammonia or through the decomposition of animal tissues upon death of the organism.

Nitrates are also transferred into the water system through runoff from the soil. Nitrates enter the groundwater. These compounds are either lost to deep sea sediments of taken up by plankton. Plankton are used as a food source for marine animals and birds in which the nitrogen is used in complex organic compounds. Decomposition of the animal tissues and animal waste return nitrates to the water and to the soil.

Nitrates in the soil may also go through denitrification and reenter the atmosphere as atmospheric nitrogen (N₂). Lightning returns the N₂ to nitrate (NO₃^-). Denitrification occurs when the soil poorly drained and low in oxygen, certain types of soil bacteria produce N₂.

The nitrogen cycle experiences many fluctuations. Nitrogen may be lost through harvesting plants, soil erosion, fire, and water runoff, but is the nitrogen fixing bacteria that continue to supply usable nitrogen into the cycle. There is a general ratio of carbon and nitrogen in the soil. Carbon is stable at 40 to 50 percent. The nitrogen level varies more, so that the range of the carbon to nitrogen ration should be between 10:1 to 30:1. If the ratio is greater, then there is a nitrogen shortage.

The aquarium is an attempt to recreate the natural marine environment on a small scale that can be observed and studied. The scale of the aquarium is very different from that of the ocean. It is very essential to produce an environment as similar to natural conditions as possible. This includes the composition of the water, the plant and animal life introduced into the plant, and also external conditions like heat and light. The composition of the water is very important but probably the most difficult aspect to establishing and maintaining a reef aquarium.

An aquarium is self-contained. It has no source of chemicals or living species besides those that are manually placed into the tank. Several chemical species serve the purpose to maintain constant conditions within the tank. These species acts as a very involved buffering system.

The bicarbonate ion plays a mayor role in buffering the system. The alkaline buffer system is in competition with a phosphate buffering system also present in the aquarium.

The carbonic acid system is an open system, meaning that the dissolved carbon dioxide may enter and leave the system. The phosphate system is a closed system. Any change in the phosphate levels must be done manually or remain at the current state.

The Henderson-Hasselbalch equation is used with buffer systems to determine the pH.

pH Ë pKa - log₁₀ [HCO₃^-]/[H₂CO₃]

This equation can be used to design and maintain buffer system around a desired pH. In the aquarium system at equilibrium, it is known that at a given pH the hydroxide ion concentration is fixed as well and the ratios between the concentration of bicarbonate, carbonic acid and carbonate.

In aquariums and tanks, water movement and the surface exchange must be assisted through the use of pumps. Oxygen is closely related to both salinity and temperature. An equation may be used to find the maximum amount of dissolved oxygen in a system according to the temperature and salinity of the water.

ln (C) = -139.34 + (1.5757_E5/T) - (6.6423_E7/T²) + 1.2438_E10/T³) - (8.6219_E11/T⁴) -

S[1.7674_E-2 - (10.754/T) + (2.1407_E3/T²)]

C = concentration of dissolved oxygen

T = temperature in Kelvin

S = salinity

An electronic probe can be used to measure the dissolved oxygen with respect to the salinity and temperature of the aquarium. A probe can also generate a three-dimensional picture of dissolved oxygen levels which may vary within an ecosystem.

In an aquarium, fish food introduces phosphates into the system. Fish waste helps contribute to the phosphate levels in an aquarium. The phosphate acts as a nutrient for algae, and plants use it for growth. As the concentration of PO₄^3- increases, so does algae growth. Plants obtain phosphates from sediment through the roots. Normal levels of phosphate are disputed. One source says that 0.09 parts per million is a normal level while another source measured phosphate levels at 0.03 parts per million.

Nitrogen compounds are very important when dealing with an aquarium. Nitrates enter the water from the animal species. These compounds, ammonia, nitrite, and nitrate, must be monitored. Increased levels are very toxic and may easily cause death to many different living species. Animal waste and decomposition of living species may be signaled by increased nitrogen levels. The appropriate nitrate levels in a fish-only tank should range from 10 to 40 ppm. Tanks that try to recreate a reef system should have nitrate levels ranging from less that 5 ppm to 10 ppm.

The aquarium in Boggs Chemistry and Biochemistry building is a 300 gallon glass tank. The ecosystem inside of the tank is a small scale reef system. A reef system was chosen because the tropical environment is easier to reproduce and maintain. There are fewer variations in the temperature and light to which the system should be exposed.

This aquarium has dimensions of 8 feet x 2 feet x 2.5 feet. The water enters the system through two pumps at either end of the aquarium. Water is also remove dfrom the tank via a gravity feed system after which the water enters a reservoir, is filtered, and in then returned to the tank again through the two pumps. There are four metal halide lights directly above the tank, producing a blue-white light. The lights are on a timer to simulate sunlight. Only one light comes on in the morning and as the day progresses more lights come on. They are turned on and off from east to west. The temperature of the tank must also be maintained. There is a lot of heat generated by the lighting system, so a fan is used to try to cool the tank. Although the fan is set to only run when the tank needs lower the temperature, the fan has never stopped turning.

The aquarium is inhabited by several types of fish from the Damsel family. There are also sea urchins and plants to balance out the coral reef. Algae is also present as a food source to the animals.

This tank is connected to the internet. Two camera provide live video of the tank to a website. One may visit the site and view the tank. The URL for this internet web page is http://science.clayton.edu/pratte/aqua/live.htm.

AgNO₃(aq) + Cl^-(aq) à AgCl(s) + NO₃^-(aq).

The potential, or cell voltage, is determined with the use of a pH meter and electrode. A silver electrode is used to detect the concentration of chloride in the solution. The overall salinity, which is the sum of all dissolved components, is calculated by the following formula:

salinity = 1.80655 [Cl- ppthou.].

Materials used included a buret, magnetic stirrer, pH meter with electrodes able to measure electrical current in mV (silver electrode), litmus paper, 250 mL beaker, a 250 mL volumetric flask, pipets, and a glass dropper. Reagents used were standardized silver nitrate (AgNO₃), concentrated nitric acid (HNO₃), and distilled water.

50 mL of water was removed from the tank. 1 mL of was placed into a 250 mL volumetric flask which was then diluted to the calibration mark using distilled water. 100 mL of the dilute solution was pipetted into a 250 mL beaker. Using a dropper, concentrated HNO₃ was added just until the solution was acidic (the acidity was checked with litmus paper), approximately two drops. An additional 2 mL of the HNO₃ was then added. The magnetic stirrer was placed into the beaker and the titration of the solution proceeded with the addition of the AgNO₃ by 0.01 mL increments. The volume in the buret and electrical current in mV from the pH meter was recorded. The data was plotted and the endpoint of the titration found before salinity could be calculated.

A blank was through the titration process to compare the results from the titration of the tank water. The procedure needs to be repeated to establish precision.

Two trials were run in the first week of testing. The very first trial had very strange results and was not taken to completion; therefore, it was redone as trial 2 during the first week. The second trial was much better, showing a clear endpoint at 3.91 mL of AgNO₃ being used. The second trial, during week 2, was found to have an endpoint at 4.00 mL of AgNO₃ used. In order to calculate the salinity of the water, the moles of AgNO₃ used during the titration must be found. The stoichiometry of the reaction between AgNO₃ and Cl^- shows a one-to-one relationship so that the number of moles of AgNO₃ equals the number of moles of Cl^- in the solution. Using the molecular weight of Cl^-, the concentration of the Cl^- ions can be found in parts per thousand. This value is then multiplied by the constant 1.80655 given from the lab manual and by 250 to make up for the 250 mL dilution of the tank water. The salinity is then reported in parts per million.

The calculated results were 30.6 ppm and 31.3 ppm respectively for the two weeks of testing. Blanks were run using distilled water and they showed a steady increase in the electrical potential after an initial jump. This steady increase verified the titration procedure and the readings from the pH meter. The blank serves to correct the data for any interaction that the pure water has during the titration. When the blank is calculated into the salinity of the solution, the slight jump would be used as the endpoint and the moles of AgNO₃ used at that point would be subtracted from the moles of AgNO₃ needed to reach the endpoint of the titration of the tank water. When the blank is subtracted from the salinity titration, the salinity concentration is calculated to be 29.8 ppm for the second trial of week 1 and 30.4 ppm for the testing during week 2.

These results are reasonable because they are similar and within range of reasonable salinity levels. Also, the salinity of the tank is normally kept below the natural salinity levels of the oceans which is around 35 ppm. The lower salinity levels are easier to keep constant, since the tank cannot renew aspects of an ecosystem. As water is added, to compensate for evaporation and water removed for testing, the salinity level may be affected.

The class results for the salinity testing were fairly inconsistent. My results were much lower than the results of the rest of the class. The range (of averages) went from about 31.5 ppm to a high of 42.7 ppm which was from the first day of testing. The average on the days which my tests were run was much higher than my results (Thursday-- week 1: class = 36.8 ppm; week 2: class = 35.2 ppm). These numbers indicate that testing on the same Thursday produced very different results. Titrations can be very sensitive procedures. Reaching the endpoint using color indicators can be very vague, but an electrochemical titration does not have an endpoint which one can miss. Errors may occur in determining where the endpoint is, once the data is graphed. The first data point for the class averages is the highest point at 42.7 ppm. This point seems unreasonable, but when the Q test is applied, the gap is not great enough to justify throwing out the point. Normal condidtions in the ocean hover around 35 ppm, and so tanks should remain a little below this level. The overall average for the quarter is high at 36.15 ppm. The standard deviation for the class results was ± 3.40. Salinity may be higher due to removal of water from the tank or increased evaporation. A ratio exists between several dissolved solvents in seawater, increased salinity levels may be in response to increased levels of other components.

CaCO₃ + H₂SO₄ (aq) à H₂CO₃ + Ca²⁺ + SO₄^2-

To calculate the alkalinity of the water, the concentration of calcium carbonate is determined. To determine the level of calcium carbonate, the second endpoint is used in order to account for all of the alkaline species Since H₂SO₄ is such a strong acid, the first hydrogen leaves the molecule very readily, creating the very acidic pH. Only the difference is needed since the reaction removing the second hydrogen actually involved work by the carbonate ions in the solution.

The moles of H₂SO₄ can be determined since the molarity of the acid is known. From the reaction between the acid and the calcium carbonate, reacting in a one-to-one ration, the milligrams of CaCO₃ can be determined. Dividing the mg of the calcium carbonate by the volume of the sample which is 0.2 L produces the alkalinity level in parts per million.

A pH meter with a pH electrode, a buret, magnetic stirrer, 200+ mL graduated cylinder, and a 250 mL beaker were used. The only additional reagent utilized was a sulfuric acid (H₂SO₄) solution standardized to 0.02 M.

200 mL was removed from the tank. The tank water was then placed into a 250 mL beaker. The beaker was placed on the magnetic stirrer with the magnetic bar. The initial pH was determined from the pH meter. Titration of the solution with the H₂SO₄ solution followed in 0.1 mL increments until the pH of the solution was 4.0. The data was potted on a graph to determine the alkalinity of the tank.

DISCUSSION:

Two titrations were run during each week of testing. In the first week of testing, the two trials generated alkalinity levels at 93.6 ppm and 94.1 ppm. These numbers are very appropiate since the normal range of alkalinity should be between 90 and 120 ppm. In the second week of testing, the alkalinity levels were increased to 99.1 ppm and 101.0 ppm. These figures are also in the appropriate alkalinity range. This alkalinity level relays that there is a fairly good buffering system in place. The carbonate species help to maintain constant tank conditions, including and especially pH. The class data shows fairly good results. Over the period of the quarter the levels seemed to increase until they suddenly dropped to below 100 ppm. Despite this change the data seems fairly good considering that all points, except one at 121, are suitable to normal tank conditions. Excluding the point at 121 due to the Q test, the average for the class was 100.71 ppm with a standard deviation of ± 6.86.

4Mn(OH)₂ + O₂ + 2H₂O à 4Mn(OH)₃

Mn(OH)₃ + 2I^- + 6H⁺ à 2Mn²⁺ + I₂ + 6H₂O

I₂ + 2S₂O₃^- à 2I^- + S₄O₆^2-.

Equipment used included the following: a buret, a magnetic stirrer, air-tight glass bottles, a 250 mL beaker, various pipets, and a glass dropper. Reagents used were manganese(II) sulfate (MnSO₄), alkali-iodide-azide, concentrated H₂SO₄, sodium thiosulfate (Na₂S2O3), and a starch indicator.

After measurement of the volume of the air-tight bottle, the bottle was filled with tank water (making sure that no bubbles were trapped in the bottle). The bottle was immediately stoppered. 1 mL of MnSO₄ and 1 mL of alkli-iodide-azide were added using a pipet. The stopper was placed back on the bottle and shaken. The solution was allowed to sit for 2 to 3 minutes. Again using a pipet, 1 mL of concentrated H₂SO₄ was added. The stopper was placed back on the bottle which was then inverted to mix the solution. Using a large pipet, 201 mL of the solution was transferred into a 250 mL beaker which was then placed on the magnetic stirrer with the magnetic bar. Titration of the solution began with the standardized Na₂S₂O₃ until the solution is of a pale yellow, or straw, color. Approximately 4 drops of the starch solution was added, so the solution changed to a dark blue hue. Titration with the Na₂S₂O₃ continued until the solution is just clear. The volume of the Na₂S₂O₃ used was recorded.

Repeat the process using water from the trough behind the tank. And repeat the entire process again.

Tests were run on both water from the tank and water from the trough. In week one the level of dissolved oxygen in the tank water was determined to be 6.16 ppm while the water from the trough was found to have a dissolved oxygen level of 6.28 ppm. In week two testing, one tank dissolved oxygen level was lower at 6.05 ppm while another was higher at 6.32 ppm. The water from the trough was lower in the dissolved oxygen levels than in the previous week with levels at 5.64 ppm and 6.27 ppm. Most of these measurements are very appropriate because the normal dissolved oxygen levels should be about 6 ppm. The class average was 6.225 ppm in the water from the tank with a standard deviation of just 0.56. My results fit into the class average very well which indicates a good chance for very accurate and precise testing.

The overall class results were fairly consistent with a range of results of only 2.1 ppm. The first week of testing did produce numbers that were higher than the rest of the quarter. This may have been due to the lack of time to develop technique. As the quarter progressed, the overall technique of each of the students should have improved producing more accurate results. The average concentration of dissolved was very similar to that of the tank. In theory, the reservoir should probably have a higher oxygen concentration.

One particular test commonly used to determine the amount of sulfate in a solution involves the addition of barium, usually in the aqueous form of BaCl₂, which contains the dissociated ions of Ba²⁺ and Cl^-. Before BaCl₂ is added, the solution is made acidic by adding hydrochloric acid (HCl). This additional reagent puts both hydrogen and chloride ions in into the solution creating a pH about 4.0. This acidic pH is the optimum environment, or characteristics, for the for precipitation of BaSO₄ to occur. The chloride ions are spectator ions and are not involved in the reaction with SO₄^2-. Once the solution is acidic, the BaCl₂ solution may be added. The barium ion reacts with SO₄^2- to produce BaSO₄, a very insoluble compound. The net ionic reaction:

Ba²⁺(aq) + SO₄^2-(aq) à BaSO₄(s).

The solubility constant of the compound is 1.1x10^-10. This K_sp is extremely small which conveys that the reaction will be shifted to the right to form precipitate. This compound falls out of solution as the precipitate of which the mass can be determined. Simple stoichiometric methods can convert the mass of BaSO₄ to the mass of SO₄^2- since the compound is formed in a one-to-one ratio.

There are slight problems that may occur with determining the amount of SO₄^2- in the solution. Peptization, the breaking up of the precipitation during the rinsing of the product may result in some product loss. Also interfering species may skew the results: Na⁺, K⁺, Li+, Ca²⁺, Al³⁺, Cr³⁺, Fe³⁺, Sr²⁺, Pb²⁺, and NO₃^- as they react with either the Ba²⁺ ion or the SO₄^2- ion. Also, the chloride ion must be sure to be removed from the product. Silver nitrate is used to test for the presence of Cl^- in the filtrate, or liquid that filtered through the ash-less filter paper. If Cl^- is present, a precipitate is formed:

Cl^-(aq) + AgNO₃(aq) à AgCl(s) [white precipitate] + NO₃^-(aq).

Warm distilled water must be rinsed through the filter until the filtrate tests negative for the presence of chloride ions. If all the chlorides ions are not removed, then the weight in the crucible after burning will be greater that than of only the BaSO₄ which will result in too high of sulfate concentration.

Materials used included the following: a desiccator, several Bunsen burners, ring stands, matches, 11 cm ash-less filter paper, funnels, a 20 mL pipet, dropper, and assorted beakers. Reagents included barium chloride (BaCl₂), silver nitrate (AgNO₃), hydrochloric acid (HCl), and distilled water.

The first step was to light two Bunsen burners, so that two crucibles may be heated and any contaminants removed. After at least 15 minutes of burning, the crucibles are placed into a desiccator to cool off the crucible. Once cooled the crucibles are measured on an electric analytic balance. The heating, cooling, and weighing were repeated until the crucibles were of constant mass.

Over 40 mL of water was removed from the tank. 20 mL of the sample was transferred into a 600 mL beaker. The sample was diluted with distilled water to an approximate volume of 250 mL. This solution was treated with HCl so that the solution is at a pH range of 4.5 to 5.0, 3 to 4 drops was added to reach the target pH. Once the target pH was reached, 2 more mL of HCl was added. This solution was then heated over another Bunsen burner until just boiling. At this point, aqueous BaCl₂ was slowly added to the solution while stirring. A precipitate formed. BaCl₂ was continued to be added until the formation of additional precipitation ceased. The solution was allowed to digest over the Bunsen burner for an hour and a half.

Once the solution had digested, the solution was removed form the heat. Distilled water was used to rinse the walls of the beaker. The filter paper was placed in a funnel and the solution was poured through the filter paper. The beaker was also rinsed with hot distilled water to ensure that all of the precipitate is removed from the beaker. The filtrate was tested for the presence of chloride by adding AgNO₃. If any precipitate formed, chloride was still present and distilled water had to be poured through the filter until the filtrate was negative for the chloride ion.

The filter paper was removed from the funned and placed into the crucible. The crucible was burned over the Bunsen burner for approximately 15 minutes. The crucible was cooled in the desiccator and was measured on the analytic balance. The heating and cooling process was repeated until the crucible was at a constant mass.

Two trials were run. The concentration of SO₄^2- can be calculated in parts per thousand.

The calculated results from the sulfate testing show that an average of 2.35 parts per thousand of sulfate was present in the water. The average by weeks is 2.4, 2.3 respectively. With three of the four trials all reporting 2.3 ppt, the results seemed very precise. The appropriate values of sulfate concentration are standard at approximately 2.75 ppt; therefore, the testing was more precise than accurate. My results were very valid although the class average was 2.149 when one weeks average of 0 was thrown out. The standard deviation of the data was only 0.22. The values of sulfur concentration were very consistent over the entire quarter. Much of this consistency can be due to the simplicity and ease of gravimetric testing.

Errors could have occurred during the determination of the sulfate concentration. Contaminated reagents would skew the results as would any dirty apparatus. Crucibles were heated and cooled so that a constant weight of the crucible could be determined; however the small amount of time for which the crucible is exposed to air does allow some moisture to be absorbed from the air. This makes the weight of the crucible a little heavier than it is actually which causes a slightly high sulfate concentration. Another error could result from not completely removing all of the chloride from the product. Hot distilled water was poured through the filter to ensure dissolution of the chloride so that it would run through the filter. The silver nitrate was used to test for the presence of chloride, and the test was negative before the filter paper was burned, but some small amounts of chloride could remain.

EDTA was used at the titrant because of its great one-to-one affinity for metal ions. The metal cations bond to the two nitrogens and the eight oxygens present in then EDTA anion.

EDTA is one of the most widely used chemicals in titrations; however, the solution pH that produces the best results varies according to the element being tested. Ca²⁺ has the most effective titrations with EDTA at a basic pH. The calcium determination needs to be at a pH near 7.5. Normal concentrations of Ca²⁺ are 400 parts per million in a marine environment.

Materials used were a buret, pipet, Erlenmeyer, Bunsen burner, and disposable glass droppers. The following reagents were also used: 0.0100 M EDTA solution, 2 M NaOH, Calcon solution, methyl orange solution, and distilled water.

50 mL of water was removed from the tank in an Erlenmeyer flask. 10 ml was pipetted into a 125 mL Erlenmeyer flask. 40 mL of distilled water was pipetted into the Erlenmeyer. 1 drop of methyl orange was added to the solution in a 125 mL Erlenmeyer flask. 2 drops of NaOH were added to make the solution a yellow color. 5 mL more NaOH was then added while stirring. 20 mL distilled water was added, and 25 drops of Calcon indicator was added to change the color of the solution to a red hue. The solution was then titrated with EDTA solution until the point at which the solution is blue.

DISCUSSION:

The titration used to determine the concentration of calcium in the sample used about 10.00 mL of the EDTA solution before the Calcon indicator indicated that the endpoint had been reached. Calculations show that the weekly concentrations of Ca²⁺ are 453 ppm and 401 ppm which average out to be 427 ppm. The normal range is closer to 400 ppm, but this shows the tests results to be fairly accurate. This is supported with the class average to be just 415.3 ppm. One value was 0.0 for one day and that data was subsequently thrown out. The overall trend of Ca2+ levels for the quarter was fairly constant which is appropiate.

The Ca²⁺ level was higher, however, than the normal concentrations in seawater. There are no large, direct effects of these higher values. Calcium is used as a nutrient by plants and by algae, which then in turn supply nutrients to animal life. The increased level of these components in the water are not at or near any toxic level. The higher levels are not serious and provide no great concern to those maintaining the tank.

Weekly fluctuations of the tank environment should not have any large effect in the concentrations of Ca²⁺ and Mg²⁺ ions. Several errors could have contributed to any inaccuracy in the testing. Contaminated reagents, inaccurate testing devices, random error, and human error could have caused problems. The titration was conducted with a color indicator. The endpoint was reached when the solution was blue. If the endpoint was just slightly off due to human error the effect would show up in inaccurate results.

Materials used were a buret, pipet, Erlenmeyer, Bunsen burner, and disposable glass droppers. The following reagents were also used: 0.0100 M EDTA solution, NH₄OH buffer, Eriochrome Blank T solution, 2 M NaOH, methyl orange solution, and distilled water.

50 mL of water was removed from the tank in an Erlenmeyer flask. 10 ml was pipetted into a 125 mL Erlenmeyer flask. 40 mL of distilled water was pipetted into the Erlenmeyer. 1 drop of methyl orange and 2 mL of ammoniacal buffer were added to the solution. The solution was not red, so no Noh was added. 20 mL more of distilled water was added and then the solution was heated over a Bunsen burner until the solution was approximately 65° C. 20 drops of Eriochrome Black T solution was added via a dropper. The solution was titrated with EDTA solution until the solution turned blue.

The titration to determine the concentration of magnesium in the tank water used approximately 60.00 mL of EDTA solution before the endpoint was reached. Calculations show that 1629 ppm and 1512 ppm of Mg²⁺ was present in the water of the tank in weeks 1 and 2 respectively. This averages out to be 1571 ppm Mg²⁺. This concentration is higher than the average of only 1350 ppm in normal seawater. The Mg²⁺ level was higher than the normal concentrations in seawater. There are no large, direct effects of these higher values. My results fit into the overall testing of the rest of the class. The class average was just 1229.5 ppmwhile the standard deviation was 95.6. All of the data points were fairly reasonable except one week which was reported to have produced a Mg²⁺ concentration of 0.0, but this is obviously incorrect and therefore was thrown out.

Some errors could have occurred during testing due to contaminated reagents, inaccurate testing devices, random error, and human error. Since the Mg²⁺ titration was conducted with a color indicator, the endpoint is somewhat vague. A slight overshoot of the titrant causes an inaccuracy in the results. Also, the solution to be titrated for the Mg²⁺ test had to be heated to between 60 to 80 ° C. In the second trial of the test, the solution went above 80 ° C. This could have had an adverse effect upon the accuracy of the testing.

H₃PO₄(aq) à H⁺(aq) + H₂PO₄^-(aq)

H₂PO₄^-(aq) à H⁺(aq) + HPO₄^2-(aq)

HPO₄^2-(aq) à H⁺(aq) + PO₄^3-(aq)

The absorbance of these free polyatomic ions are then measured in the spectrophotometer. Beer's Law shows that the concentration of the species is directly proportional to the absorbance of the solution.

Materials used were a filter, pipets, 300 mL Erlenmeyer flasks, 250 mL volumetric flasks, disposable droppers, and a spectrophotometer. Nanopure water, phenolphthalein indicator, strong acid solution, ammonium molybdate ((NH₄)₆Mo₇O₂₄), stannous chloride (SnCl₂), and standard phosphate solution were the reagents used.

Standard phosphate solutions must be prepared. 10 mL of the standard phosphate solution was transferred via pipet into a 250 mL volumetric flask to form standard #1. The solution was then diluted to the calibration mark using nanopure water. The same process is repeated for standards #2 and #3 by adding 2 mL and 1 mL respectively.

To begin preparing the solutions to be tested, 250 mL of water was removed from the tank and run through a filter. The filter was presoaked in nanopure water. Two 50 mL amounts of the filtered sample was transferred into two 300 mL Erlenmeyer flasks, A and B. 50 mL of standard #1 was pipetted into flask C. 50 mL of standards #2 and #3 are transferred into their own Erlenmeyer flasks, D and E. 50 mL of artificial seawater is placed into flask F while 50 mL of nanopure water is placed into flask G. 50 mL of nanopure was placed into each of the seven flasks. One drop of phenolphthalein indicator is added to each of the flasks. None of the solutions turned pink; therefore, no acid was added. Next, 4 mL of the molybdate reagent and 10 drops of SnCl₂ were added. The solutions were thoroughly mixed. Once the solutions had time to settle for between 10 and 12 minutes, the absorbance of the solutions were determined through the use of a spectrophotometer. To zero the machine, a cuvette of nanopure water was used. To determine the maximum, or wavelength of peak absorbance, a cuvette with the solution from flask C which contained the standard #1, was used. The spectrophotometer measured the absorbance at a range of wavelengths to determine the wavelength at which the absorbance is the greatest. The computer generated a graph to find the peak wavelength of peak absorbance. On a separate spectrophotometer, this wavelength was used to determine the absorbance of the other solutions. The absorbance of each solution was used to calculate the concentration of PO₄^3-~P in the water.

Material and supplies needed were Teflon® bombs as the digestion vessels, microwave, 100 mL volumetric flasks, pipets, droppers, spectrophotometer, and cuvettes. Two reagents used were the filtered tank sample and standard phosphate solutions that were made in part II. Other reagents included phenolpthalein, 6M NaOH, nanopure water, molybdate reagent ((NH₄)₆Mo₇O₂₄), and stannous chloride (SnCl₂).

50 mL of two tank samples, the three phosphate standards, artificial seawater, and nanopure were put into their respective Teflon® bombs. One drop of phenolphthalein indicator was added to each vessel to see if any acid needed to be added, but the solutions were not pink; therefore, no acid was added. The seals were placed on the Teflon® bombs. The tray of vessels was then placed into the microwave. The vent hoses were connected to control the pressure inside each of the bombs, and the microwave was programmed to run a waste water program. The program ran for about 20 minutes at 95 percent power, and once the pressure inside each of the vessels returned to a pressure below 5 psi, the vessels were removed. The solutions were transferred into 100 mL volumetric flasks. 6 M NaOH was added just until the solutions were of a light, faint pink hue. The solutions were diluted up to the calibration marks with nanopure water. 4 mL of molybdate reagent and 10 drops of SnCl₂ were added to each flask with thorough mixing. The solutions were allowed to sit for 10 minutes. At this point a cuvette of nanopure water was used to zero a spectrophotometer. The same cuvette was rinsed and filled with the solution of standard #1. This cuvette was placed into the spectrophotometer so that the peak wavelength could be found. The absorbance was measured at a range of wavelemgths, from 400 to 850 nm. The wavelength of maximum absorbance was used to measure the absorbance of the other solutions. The same cuvette was rinsed and filled with each of the other six solutions so that the absorbance could be determined. The absorbencies for the different solutions allowed for the concentration of PO₄^3-~P to be determined.

In order to find the concentration of phosphates in the sample tank water, solutions were made and then compared to prepared standards in relation to their absorbency in a spectrophotometer. Graphs were produced by plotting the concentration of the standards versus the absorbance at a optimum wavelength. An equation was generated through the origin. This equation was used to find the concentration of the unknowns. The absorbance data of the standards was very precise. The equations of the lines were similar with nearly equivalent slopes.

To find the concentrations of the unknowns, the absorbance of the blank, a cuvette of nanopure water, was subtracted from the absorbance of the tank sample and the absorbance of the artificial seawater. For an unknown reason, the absorbance of all of the blanks was higher than the absorbance of the samples being tested. This could be the result of contamination, malfunctioning equipment since the same spectrophotometer was used for all of the measurements besides the determination of the wavelength of the maximum absorbance. Completion of the calculations result in negative concentrations of phosphate. This data can also be used to infer phosphate value of 0 which is very likely, even for a healthy aquarium. The normal values of phosphate should be around 0.07 to 0.09 parts per million.

Although the numbers produced negative values for the concentration of phosphates in the water, comparison of the results of the tank water samples to the seawater sample provided for some interesting conclusions to be drawn. The absorbance of the seawater was very similar to that of the two tank samples. The seawater absorbance was only a little higher than the tank samples. In week 1, the absorbance of the seawater was 0.030 while the average absorbance of the tank samples was 0.024. In week 2, the seawater was just 0.023 while the tank samples were 0.020. These similarities suggest that the error is more likely in the measurement of the absorbance of the blank and not in the measurement of all of the other samples.

To try to obtain some kind of phosphate concentration numbers, the blank may be assumed to be near negligible. In this case, the concentrations are found using the same derived equations. These calculations produce an average experimental concentration of 0.0130 parts per million. This number is very low compared to the disputed normal range of 0.03 or 0.09 parts per million, but you must consider that the absorbance was not corrected by the use of the blank. The results from week two, without subtracting the blank are 0.0121 ppm.

Most of the results of the rest of the class over the period of the quarter were also very close if not 0. Ppm. Four individual results were actually above 1 ppm. The highest value found was 12 ppm. This data is obviously an error. For the most part, the tank was very constant in phosphate levels as should be expected since any phosphates in the water act as a buffering system to stabilize the aquarium and provide nutrients for plant growth.

Many of the same procedures followed in measuring acid hydrolyzable orthophosphates are similar to those in the determination of reactive orthophosphate. However, in measuring acid hydrolyzable orthophosphate, the samples were placed in Teflon® bombs and heated and pressurized to help dissociate some of the phosphate compounds. In the phosphate determination without the microwave digestion, some of the phosphate compounds may not have been included in the absorbance measurement; therefore, the phosphate concentration should be higher in the microwave digestion testing.

The absorbance of the samples must be corrected by subtracting out the blank, but due to an error, the blank is higher than the other measurements. This produces a negative absorbance and, therefore, a negative concentration of phosphate. The concentrations were extremely low with an average value of -0.034 ppm PO₄^3-~P. This most likely indicates that the concentration of phosphates in close to if not 0.0 ppm. If the blank is considered to be negligible, just to obtain some type of phosphate concentration although high due to the lack of correction of the blank, the results acquired are consistent within each week of testing but not between the two weeks. The first week produced results with higher phosphate levels. The average of the tank water is 0.0268 parts per million PO₄^3-~P. This result is close to the experimental value of the artificial seawater. The idea that the PO₄^3-~P levels should be higher that the phosphate determination of that same week holds true, but the numbers are still low in comparison to the normal concentrations. The results of the microwave digestion of week 2 indicate PO₄^3-~P levels of 0.0097 ppm. This average is much lower than any of the other tests or trials.

Errors may have occurred in lab occur contamination of reagents, malfunctioning equipment including the spectrophotometer or microwave, human error, or even random error, but my data was similar to the rest of the class. Besides two very strange results near 40 and 14 ppm, all of the results are very close to 0.0 ppm. This consistency between trials, time, and people indicate that the levels are relatively constant and the data both accurate and precise.

The determination of nitrite concentration used spectrophotometry to measure the absorbance of the solution. Beer's law states that absorption is directly proportional to concentration. Standards are used to generate a linear equation and determine the wavelength of maximum absorbance. The color reagent serves provide a color basis for the absorption. A greater concentration of nitrite coincides with an increased intensity of the color shown by the color reagent.

The final experiment which is used to determine the concentration of nitrate uses a copper-cadmium column. The copper-cadmium column is used to reduce the nitrate ion to nitrite as the solution is drained through the column:

2NO₃^- à 2NO₂^- + O₂.

The concnetration of nitrite is then determined through absorption sppectrophotometry. This results in an actual determination of nitrate plus nitrite.

Materials needed are an ammonia selective electrode and pH meter, magnetic stirrer, beakers, pipets, and volumetric flasks including 1 1-L, and 2 100-mL. Reagents that were used were ammonium chloride (NH4Cl) stock solution, 10 N sodium hydroxide (NaOH), artificial seawater, and distilled water.

The first step was to prepare ammonium standard solutions of approximately 1.0 ppm, 0.5 ppm and 0.1 ppm. The stock solution was prepared to 3188.8 ppm. The is a conversion factor for 3188.8 ppm stock solution: 1 mL = 1.22 ppm NH₄; therefore, 1 mL of the stock solution was transferred via pipet to a 1 L volumetric flask. The flask was then filled to the calibration line using artificial seawater producing a standard concentration of 1.22 ppm. To make a more dilute solution, 5 mL of the 1.22 ppm solution was pipetted into a 100 mL volumetric flask. This flask was then filled to the calibration mark with artificial seawater producing a 0.61 ppm solution. Finally, 1 mL of the 1.22 ppm solution was pipetted into another 100 mL volumetric flask and diluted producing a 0.122 ppm solution. Although the solutions are not exactly 1.0 ppm, 0.5 ppm, and 0.1 ppm, as long as the concentrations are known and cover range of concentrations, the standards will be suitable for the testing.

350 mL of water was removed from the tank in a graduated cylinder. 100 mL of this water was pipetted into a 200 mL beaker with a magnetic stirring bar. The beaker was placed on the magnetic stirrer, at a slow setting. 1 mL of 10 N NaOH was pipetted into the beaker. The initial millivolt reading was recorded was the reading was recorded in thirty second intervals until the solution was stable. This took less than 2 and a half minutes. Three trial of the tank water were run. The three standards were also run through the procedure as was an artificial seawater blank.

Materials used include the following: 100 mL volumetric flasks, 50 mL Erlenmeyer flasks, cuvettes, spectrophotometer. Nitrite stock solution, color reagent, distilled water, and artificial seawater were used. Standard solutions were made from the nitrite stock solution.

100 mL of water was removed from the tank. 25 mL was transferred to a 100 mL volumetric flask via a pipet. This solution was diluted to the calibration line using distilled water. The same process was repeated again with another sample of tank water and also with the three standards, and a artificial seawater blank. 50 mL of each solution was then transferred into 50 mL Erlenmeyer flasks. 2 mL of the color reagent was added to the solution and mixed well. After ten minutes passed, the absorbance of the 1 ppm nitrite standard was determined over a wavelength range from 400 nm to 850 nm. The wavelength of maximum absorption (l = 541.0 nm) was then used to acquire an absorption reading for the rest of the solutions.

Materials used were a buret with a column of copper-cadmium granules, (1 L, 100 mL, and 50 mL) volumetric flasks, pipet, funnel, and (125 mL, 50 mL) Erlenmeyer flasks.

Reagents used were the cadmium column, nitrate stock solution concentrated ammonium chloride-EDTA solution (NH₄Cl-EDTA), distilled water, and a coloring agent.

Nitrite standards had to be prepared from the KNO₃ solution. The conversion factor from the KNO₃ solution to just the NO₃^- ion is 1 mL of KNO₃ solution = 100 m g NO3-. Since 100 m g equals 0.1 mg, 10 mL of the KNO₃ solution was diluted to 1 L in a 1 L volumetric flask, diluted with artificial seawater. The 0.5 ppm solution was made by pipetting 50 mL of the 1.0 ppm solution into a 100 volumetric flask and diluting to volume with artificial seawater. Similarly, the 0.1 ppm standard solution was made by transferring 10 mL of 1.0 ppm solution into a 100 mL volumetric flask.

Seven solutions were carried through the procedures. 10 mL of each solution was placed into separate 50 mL volumetric flasks. The solution was diluted to the calibration mark with distilled water. 25 mL of this diluted solution was pipetted into an 125 mL Erlenmeyer flask. 75 mL of concentrated NH₄Cl-EDTA solution was added to the Erlenmeyer which was then mixed thoroughly. A buret had about 10 cm of the copper-cadmium granules and some solution (NH₄Cl-EDTA). The buret was drained until approximately 1.5 cm of fluid remained above the granules. At this point, about 25 mL of the 1 ppm standard was poured into the buret. The solution was drained at a rate of 7 mL per minute until the level of the liquid was about 1.5 cm above the copper-cadmium granules. The solution that was drained through the column was discarded in a waste beaker. The remaining 1 ppm solution was added to the buret and drained through the column. 50 mL of this drained solution was collected in an 50 mL Erlenmeyer. 2 mL of the color reagent was added and mixed. There was no color change. The standard of 0.5 ppm concentration was then drained through the column, discarding the first 25 mL. The 50 mL collected showed no color change with the addition of the color reagent; therefore, the testing was not completed.

To calculate the concentration of ammonia in the water. The electrical potential, recorded once the system stabilized at equilibrium, was plotted on a graph versus the negative log of the concentration for each of the three standards. The use of the log function serves to create a linear graph rather than either exponential growth or decay. The data should produce a straight line with a negative slope since the higher concentrations of ammonia would produce a greater concentration of ions to create the higher electrical potential; however, this is not the case with the acquired data. The millivolt reading of 19 for the 1.22 ppm standard solution is very far from the other data and from the expected value. The graph does not resemble a linear function. Since this point is determined to be the incorrect data, it can be thrown out and a line can be generated using the two points from the 0.61 ppm and 0.122 ppm standards. This line is not as accurate, since it uses only two points, but may still be used to find the concentration of ammonia. The equation for the line is

y = -10.015x + 33.15

where 'y' is the mV reading and 'x' is the negative log of the concentration of the standards in parts per million.

Three trials were run of the water from the tank. The mV readings were 32, 26, and 25 respectively, producing calculated concentrations of ammonia to be 0.7677 ppm, 0.1932 ppm, and 0.1535 ppm. These numbers produce and average ammonia concentration of 0.3715 ppm. This value seems reasonable since the average value of the ammonia concentration should be approximately 0.1 ppm. The first trial, although, seems to be high in comparison to the other two trials. If the first trial is thrown out, the average of the two other readings gives a ammonia concentration of only 0.1734 ppm. This data may be low but is still reasonable since low levels of ammonia are desired.

The data produced in these particular tests are reasonable when the class data is taken into consideration. The beginning of the quarter produced higher ammnia levels closer to 2 ppm which are high. As the quarter progressed, the levels came down to levels below 1 ppm. This data is very reasonable.

Several errors could have occurred in this lab. The electrical potential of the first standard was incorrect. This could have been due to a problem with the pH meter/electrode, but this in unlikely since the following data obtained was fairly good. There could have been contamination or human error. One fact to note was that the concentration of the first standard was 1.22 ppm, above the proposed 1 ppm solution. This increased concentration could have been out of calibration range for the electrode, but this in unlikely since even the most ammonia sensitive electrodes measure up to 10 ppm.

Absorption spectrophotometry was used to obtain the concentration of the nitrite ion (NO₂^-) in the water of the fish tank. Beer's Law was used to justify the use of a linear graph of absorbance versus concentration. The following linear equation was used to determine the concentration of nitrite in the water

y = 0.47x

with 'y' representing the absorbance and 'x' equaling the concentration in parts per million. Using this equation, the nitrite concentration was found to be 0.0085 ppm for both trials of the tank water and 0.0128 ppm for the artificial seawater. These results are very good since the lower the levels of nitrite in the water, the better. The normal levels of nitrite are very close to zero, close to the results obtained.

The results discussed above used a Milton spectrophotometer. The same solutions were also tested on the older and non-digital spectrophotometers in the lab using the same peak wavelength obtained earlier from the scanning Milton spectrophotometer. The readings from these older machines were close to those from the previous testing. The equation generated from the graph of the standards was

y = 0.4594x.

Using this second equation the nitrite concentration was found to be 0.000 ppm for both trials of the tank water and 0.0174 ppm for the artificial seawater. Some of the difference in the results between the two spectrophotometers could have been due to the use of plastic versus glass curvettes, but the similarity of the values from the two spectrophotometers rule out any large error on the part of the machines.

The class data was not very similar to the data obtained from my tests which had levels very close to 0.0 ppm. The average of the other data was 1.1 ppm. This may be high, but it is still reasonable. Lower levels of nitrite are desired in a tank environment.

Even though similar results were found using two spectrophotometers, errors still could have occurred during the lab to skew the results. Contamination or human error is always a probability in lab. Contamination of the reagents could basically translate the data either up or down producing a high or low concentration of nitrite.

The copper-cadmium column was used to reduce the nitrate ion to a nitrite ion of which the concentration could be determined using spectrophotometry. When the standard solutions were run through the copper-cadmium column, the ions were not reduced, because at the addition of the coloring reagent, there was no visible change in color. A strong pink color should have appeared with the addition of the reagent. Another standard was run through and tested with the coloring reagent but again there was no apparent color change. If the standards did not produce a color change with their high levels of nitrate, no equation could be produced to determine the concentration of the samples, so the testing was not carried out.

This was not a surprise that the column did not work, since it has been very rare that it has worked at any time this quarter. When using a copper-cadmium column to reduce nitrate, the column must be activated at regular intervals, and it must not be allowed to be drained dry or the column would be reduced in a reaction with the air. The error is most likely with the column but it may also be with the color reagent.

Many people in industry have trouble getting results with copper-cadmium columns. Many people are trying to find different and more reliable ways to measure nitrate levels. One company, The Nitrate Elimination Company, Inc., is currently developing a BioSensor ®. This electrical device uses the enzyme nitrate reductase (NaR) to convert nitrate to nitrite. This enzyme is paired with a natural coenzyme NADH. This coenzyme, naturally derived from yeast, acts as the electron donor. These reduced samples may then be tested with dyes to detect the concentration of nitrite. This method might should be considered as an alternative for the temperamental copper-cadmium column.

The entire class and very negative results when attempting to find nitrate concentrations using the copper-cadmium column. Results that were obtained showed higher levels ranging from 0.3 pmm to 2.6 ppm. The testing was not proven to be very accurate or reliable over the course of the quarter.

The role of sulfate in the marine aquatic system is not completely understood. The effects of the lower levels of sulfate are not directly known either. It is known that the tank which was examined, runs at a lower salinity concentration than the natural marine ecosystem. Salinity is a measurement of all dissolved ions in the water, one of which is SO₄^2-. Lower salinity levels lead to lower SO₄^2- levels, so that the low results are very reasonable. The constancy of the results from one week to the next is also another point that reinforces the quality of the testing. Observations of the tank showed no drastic change or large problems with the water quality. All of the species seemed to be living well. The low sulfate levels probably do not have any adverse effect, or at least any that can be seen with the naked eye and without more extensive testing of the life within the tank.

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