The concentration of soluble phosphates in natural water samples (orthophosphate and condensed forms) will be determined using standard methods adapted to small scale techniques.
This lab is designed for first year chemistry students to use at the end of the year, an AP chemistry class or for an environmental science class with students who have a basic understanding of chemistry. Students will create phosphate standards and determine concentrations using small scale techniques. Matching of colors for colorimetric analysis is done by simple standard comparison or two photometric methods. An inexpensive permanent color chart may be created to use for further soluble phosphate determinations.
Approximately 50 minutes.
Reagents prepared for this lab exercise will be in excess, enough to serve over 100 students. Preparation in bulk is advised following these directions to eliminate weighing errors. Dilution procedures are designed to minimize potential sources of analytical error.
Use goggles and plastic aprons for the entire laboratory. Do not use phosphate containing cleaners on any sample containers or labware as residue will contaminate analyses. A dilute acid bath for soaking and rinsing of labware is recommended. Several acids and potentially toxic chemicals are used in this laboratory exercise. Use extreme care in handling acids and other chemical solutions. Follow lab directions exactly unless otherwise advised.
Working standards prepared by students:
|Tube #1||10 ppm||Tube #2||5 ppm|
|Tube #3||2.5 ppm||Tube #4||1 ppm|
|Tube #5||.5 ppm||Tube #6||.2 ppm|
|Tube #7||.1 ppm||Tube #8||0 ppm|
Sample Calculation of Phosphate Working Standard Values:
The phosphorus concentration is determined using KH2PO4 as a primary standard determined by mass. The atomic mass of the compound is 174.18 g/mol. The atomic weight of P is 30.97 g/mol. To determine the amount of KH2PO4 for weighing to have approximately one gram of P, multiply the atomic weight of the standard (174.18 g KH2PO4 per 1 mol KH2PO4) times (1 mol KH2PO4 per 1 mol P) times (1 mol P per 30.97g P) times (1.00 g P/1 liter of solution). This gives 5.62g KH2PO4 per 1 liter of solution that needs to be prepared for weighing. The true weighed value does not need to be this exact amount but does need to be reported to the correct number of significant figures to determine the concentrations made by successive dilutions.
The first concentration of P is made by by dissolving 5.62g of the primary standard/liter of solution and gives 1g P per 1000g of solution or 1 part per thousand which is equivalent to 1000 ppm.
Take 10mL of this stock solution and dilute to 1000mL (10mL of 1000ppm solution) times 1/1000mL = 10microgram per 1000mL or 10ppm.
This primary dilution from the 1000ppm stock solution is the parent for all further dilutions.
Take 25mL of the standard solution and dilute to 50mL.
25mL(10 ppm) X 1/50 = 5ppm
Take 10mL of the standard solution and dilute to 100mL.
10mL(10ppm) X 1/100 = 1ppm
Take 2mL of standard solution and dilute to 100mL.
2mL (10ppm) X 1/100 = 0.2ppm
Take 1mL of standard solution and dilute to 100mL.
1mL (10ppm) X 1/100 = 0.1ppm
This logic, while done in milliliter units, parallels the lab procedure which is written in drop ratios.
Obtain a piece of thin light blue mylar. The plastic strip should be close to the color of lightest standard. A permanent colorimetric standard may be made by cutting the blue mylar into strips of varying lengths, offset by a couple of centimeters. Once 5-6 strips have been cut in successively shorter lengths (the same offset dimension), the strips may be stacked like cards in a playing deck and stapled (like a paper pad) on one end. Staples may be placed in the long dimension sides to avoid "fluttering" of the strips. The entire assembly may be heat sealed in a small plastic sleeve for protection. Compare the successively darker blue colors with your known colorimetric solution standards and mark the concentration for each color section. This is a new "secondary" standard from which to perform colorimetric analyses. Protect the color strip, once made, from exposure to bright light or scratching.
Prepare standards in the same manner as method #1a. Allow color to develop. Follow specific directions for the microspectrophotometer as provided (Stelter, Tunt, and Lieneman (1990), this volume. Use tube #8 (the "blank") to read the "zero" resistance value.
If a blocktronic, Spectronic 20, or other commercial spectrophotometer is available, prepare standards in the sample manner as method #1a. Allow color to develop. Follow specific directions for the determination as instructed. Set the "no light" condition with the shutter closed. Use tube #8 to set the full scale absorbance reading.
Using any of the methods, appropriate readings (qualitative or quantitative) must be done on all samples. Record all data generated and take appropriate numbers of readings to insure that the values obtained are valid. Create a sequence (qualitative analysis) or a calibration graph (quantitative analysis) using the concentration values as the independent (x-axis) variable and the "response" values as the dependent (y-axis) variable. Report your results in tabular and/or graphical form as indicated.
Water samples for phosphorus determinations should be collected in glass bottles that have been rinsed with an acid wash solution, rinsed with distilled water, and, just prior to sampling, rinsed several times with the water being collected. A record should be made of the location of sample collection by recording exact, location, time, temperature, and name of the person who performed the sampling. To preserve samples, filter immediately and freeze. If samples are to be kept for long periods of time, add 40 mg HgCl2/liter to kill living organisms. If not frozen, samples may be kept 48 hours if refrigerated. Do not store in plastic bottles because phosphates, at very low concentration levels common in natural waters, may be adsorbed onto the walls of plastic bottles.
Analysis of Water Samples
Orthophosphates and condensed phosphates are highly soluble in water. The orthophosphate ion (PO43-) has the P atom centrally located, bonded to oxygen atoms which are located at the corners of a tetrahedron. Condensed phosphates, the polyphosphates, and metaphosphates are formed by condensation of two or more orthophosphate groups and have a P-O-P linkage. The polyphosphates are linear; metaphosphates are cyclic. Condensed phosphates are the most abundant form of phosphates in natural waters.
Phosphates enter lakes, ponds, rivers, estuaries, and the ocean from various primary sources such as inorganic fertilizers, wastewater treatment from municipal sources, runoff from feed lots, soaps and detergents, and industrial processes. Certain types of detergents can introduce a high concentration of phosphate ions into bodies of water. In detergents, tripolyphosphates are used to stabilize dirt particles and complex Ca2+ and Mg2+ to prevent combining with the detergent molecule resulting in superior cleaning ability. However, when the soluble detergents are rinsed away, the resulting wash water has high concentrations of phosphates which act as environmental "nutrients."
Phosphates and nitrates are often limiting factors for populations of aquatic plants, algae, and other vegetation. Phosphates are the primary limiting factor in fresh water plant and algal growth. Nitrates are believed to be the primary limiting factor in control of marine plant growth. However, the synergistic effects of essential trace elements must be considered in both environments. Large amounts of phosphates in a body of water tend to stimulate growth of algae and aquatic plants. If levels become too high, plant growth can accelerate resulting in rivers and lakes with dense growth of algae and plants. This is referred to as an "algal bloom". In 1980, 48% of the lakes and rivers in the United States exceeded the 0.1 mg/liter phosphate standard to prevent the growth of nuisance plant life. A level of .025mg/liter will result in an increase in growth of plant life.
Accelerated and uncontrolled plant growth impairs fishing, boating, swimming and other recreation uses of natural waters as well as the natural ecological balance in the particular environment. In the Fall in temperate freshwater environments seasonally active plants die and other vegetative growth decreases substantially. Dead and decaying plant materials are decomposed by aerobic bacteria which, depending on growth rates of populations, can deplete the oxygen content in the water. This oxygen depletion or minimum causes respiring aquatic organisms to die. Anaerobic bacteria assume the role of decomposition under conditions of very low oxygen concentrations. In the breakdown of the organic material by anaerobes, toxic byproducts are often produced.
Natural succession in aquatic environments due to increased nutrient loading is referred to as "natural eutrophication." In this process, a clear shallow lake becomes cloudy and ultimately begins to fill in, becoming a swamp. Eventually, as sedimentation continues, the water body becomes dry land over a period of hundreds of years. When this process is accelerated by human influence, primarily by the introduction of excessvie phosphates into the water, it is referred to as "cultural eutrophication."
Reactive phosphorus is largely in the orthophosphate or condensed form, which is soluble in water. These phosphates respond to the colorimetric detection method explained and determined by operational definition. Particulate and filterable phosphorus may be converted to soluble orthophosphate or condensed form by acid hydrolysis at boiling temperatures. Oxidative digestion for organically bound phosphate may also be determined by some method of chemical or physical sample pretreatment prior to analysis. These "forms" of phosphate in the natural environment are also "operationally defined" by the means of analytical determination. Only the naturally occuring soluble form will be will determined in this lab.
Orthophosphate or condensed phosphates react with molybdate in strong acid solution to produce heterpolymolybdate. Other forms of phosphates do not react under these defined conditions. The heterpolymolybdate that is formed is phosphomolybdic acid: H3(Mo3O10)4, a yellow compound with maximum absorption at 350nm. This is too low for absorptiometric method of water analysis to be used for determinations. If the molybdenum is reduced by the ascorbic acid to a molybdenum blue compound with uncertain concentration ratio of Mo(IV) and Mo(VI), a colloid with maximum absorption near 700nm is produced. With additions of trivalent antimony made to the acid molybdate reagent, modified heterpoly acids are formed and reduced to a modified molybdenum blue compound, the composition of which is generally unknown. This is the compound which is used colorimetrically for analytical work.
In this procedure, ammonium molybdate and potassium antimonyl tartrate react in acid medium with orthophosphate to form a heterpoly acid, phosphomolybdic acid, that is reduced to intensely colored molybdenum blue by ascorbic acid. Different intensities of blue for different concentrations of PO43- are produced allowing a reproducible determinative method to be applied.
The small volumes of reagents and materials used in this activity may be disposed of with large dilutions by flushing solutions down the drain. Excess quantities of metal containing reagents as bulk solutions should be stored and/or disposed of as identified metal ion waste.
Method adapted from Standard Methods for the Examination of Water and Wastewaterfor small scale lab activities involving microspectophotometry and a permanent colorimetric system.
American Public Health Association (1988). Standard Methods For the Examination of Water and Wastewater, 15th ed. Washington, D.C.: APHA, pp. 409-421.
Halmann, M; ed. (1972). Analytical Chemistry of Phosphorus Compounds, New York: John Wiley & Sons, p. 248.
Snoeyink, Vernon L. and Jenkins, David (1982). Water Chemistry, New York: John Wiley and Sons, p. 299.