Introduction
For this experiment the objective was to utilize spectrophotometry and the use of Beer’s
law while determining an unknown phosphate concentration of a water sample based on a
calibration graph for phosphate content. All bodies of water contain a level of organic phosphate,
though when phosphate levels rise they can have detrimental effects on aquatic life and plant
growth. A natural level of phosphate ranges from .0005 to 0.05 mg/L any phosphate level greater
than 0.1 mg/L is considered detrimental to water quality and aquatic life3. Phosphate can come
from natural sources like a toilet flush or is present in organic or polyphosphate form in
fertilizers and pesticides. Dissolved phosphorus (usually present as PO43- ) is in one of three
forms: orthophosphates, condensed phosphates (polyphosphates) . Orthophosphate is readily
available for microbial uptake which is why it impacts water life so heavily. Eutrophication is the
effect of excessive amounts of nutrients in a body of water, likely due to runoff from the
surrounding land/soil. Eutrophication effects are: dense plant growth and death of aquatic life
from lack of oxygen in the water. An example of extreme eutrophication is currently in the Gulf
Of Mexico. There is a “dead zone” in the Gulf of Mexico due to excessive nutrients from
agriculture near the Mississippi River which have ran off into the Gulf2. Phosphorus promotes
photosynthesis, therefore when phosphorus is in abundance, there is overgrowth of plants/algae
in bodies of water. Overgrowth of plants/algae depletes oxygen in the freshwater source. The
nutrient enrichment of the Gulf has created an oxygen-deprived dead zone. Dissolved phosphate
concentrations can be determined through the use of spectrophotometry to measure absorbance
of the sample. Spectrophotometry was used to obtain measurements of absorbances at particular
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wavelengths. The spectrophotometer separates lights into wavelengths and measures the intensity
of light at each wavelength. A detector converts the amount of light passed through the sample
into an electrical signal which is converted into a reported absorbency of the sample. The
concentration of an unknown was calculated using a calibration graph, measured absorbency, and
Beer’s Law. Dilutions of water samples were done using the equation C1V1=C2V2 was
manipulated to (ppm desired/ppm stock solution)(100 mL)= volume of stock solution. A
calibration curve was created using the known concentrations of phosphates from diluted water
samples. The calibration curve gave a R-squared value and an equation so a concentration of the
unknown could be determined using the equation of the standard curve. Beer's Law was used
because it directly relates the amount of absorbing species present, amount of solution the light
passes through, and absorptivity of a solution. The formula of Beer's Law used was A= ϵ lC.
Where A is absorbency, ϵ is the absorptivity constant, l is the path length, and C is the
concentration of the absorbing species.
Methods
To begin, 6 empty clean beakers were obtained and one of each were labeled with .5, 1.,
2.0, 3.0, 4.0, or 5.0, indicating the concentration of 20 ppm solution of phosphate ion in the
beaker. A 50 mL buret was filled to the 20 mL line with phosphate. For the first dilution, 2.5 mL
of 20 ppm solution of phosphate ion was measured out from a buret and into a 100 mL
volumetric flask, then the flask was filled to 100 mL with distilled water (~97.5 mL water) and
transferred to the beaker marked with concentration “.5”. For the second dilution, 5 mL of 20
ppm solution of phosphate ion was measured out from a buret and into a 100 mL volumetric
flask, then the flask was filled to 100 mL with distilled water (~95 mL water), and transferred to
the beaker marked with concentration “1.00”. For the third dilution, 10 mL of 20 ppm solution of
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phosphate ion was measured out from a buret and into a 100 mL volumetric flask, then the flask
was filled to 100 mL with distilled water (~90 mL water),and transferred to the beaker marked
with concentration “2.00”. For the fourth dilution, 15 mL of 20 ppm solution of phosphate ion
was measured out from a buret and into a 100 mL volumetric flask, then the flask was filled to
100 mL with distilled water (~85 mL water),and transferred to the beaker marked with
concentration “3.00”. For the fifth dilution,20 mL of 20 ppm solution of phosphate ion was
measured out from a buret and into a 100 mL volumetric flask, then the flask was filled to 100
mL with distilled water (~80 mL water), and transferred to the beaker marked with concentration
“4.00”. For the sixth dilution, 25 mL of 20 ppm solution of phosphate ion was measured out
from a buret and into a 100 mL volumetric flask, then the flask was filled to 100 mL with
distilled water (~75 mL water), and transferred to the beaker marked with concentration “5.00”.
To find the absorbency of the phosphate, a spectrophotometer was used. A “blank” was mde used
distilled water to ensure all samples were calibrated from the same starting point. For each
beaker, a 20 mL pipette was used to pipette 20 mL of the diluted sample into a Erlenmeyer flask.
Then using a 1 mL pipet 1 mL of molybdate solution was added to the sample and was mixed.
Then, 2 drops of Sn(II)Cl2 was added to the water sample and was mixed. The time was noted
after all components were added to the sample and after 6 minutes the absorbency was calculated
using the spectrophotometer. This same process was used for all 6 samples of the water, blanking
the spectrophotometer before each reading. Each absorbency, with its correlating concentration
of phosphate was recorded in Excel and a scatter plot of the data was created. The data was used
to obtain a trendline and an equation of the best fit line. After the standard curve was made, the
chosen local water sample, Showalter Fountain, could be analyzed. Three separate 20 ml
portions of the unknown were made along with 1 ml molybdate and 2 drops of Sn(II)Cl2. . These
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samples were put into the spectrophotometer after 6 minutes at 650 nm. Using the best fit line
equation, the concentration of phosphate in the unknown water samples were calculated.
Results
Figure 1: The following figure is the standard curve or calibration curve made from the diluted
phosphate samples in Part 1 of the experiment. The calibration curve generated a best fit line
equation and an R-squared value.
Standard Phosphate
Concentration (ppm)
Volume of 20.0 ppm
phosphate solution
required to make 100.00
Absorbance