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					Melissa Friscia & Matt Cunningham

  Real World Project: Determination of Iron in Blood Using Iodometric Titration
Objective: Using the Iodometric tactics we used in Quantitative Analytical Chemistry to determine the
amount of iron in a fresh blood sample, and compare the values obtained to a six month old sample of
blood to determine if iron depletes out of blood after six months.

Background:

Iron is the main component of hemoglobin, which is the Oxygen-transporting substance for the body
(via blood). This transport is accomplished by way of a reversible reaction between Oxygen and the
hemoglobin. A lack of Iron in the blood can lead to fatigue, lack of appetite, increased risk of infection,
shortness of breath, and anemia. Anemia is accompanied by decreased red cell production, acute blood
loss, and an increase in hemolysis (destruction of red cells).

Procedure:

Standardizing Thiosulfate:

1. A mass of 1.0912 grams of KIO3 was massed out and diluted to 100 mL in a volumetric flask

2. 10 mL aliquots were taken of the KIO3 solution to use each titration (per Erlenmeyer flask)

3. Approximately two grams of KI were massed, and added to each Erlenmeyer flask; this caused the
solution to turn brown

4. The flasks were then titrated with Thiosulfate solution

5. When the solution in the flask turned to about the color of white wine, a 2 mL sample of 1% starch
solution was added, and this turned the solution a dark purple bluish color

6. The solution of the flask was titrated to a clear endpoint

7. Then calculate the molarity of the Thiosulfate solution

Creating an Iron Standard (FeCl3 1000 ppm)

1. 0.1 grams of FeCl3 were weighed out

2. The 0.1 grams of FeCl3 were diluted to 100 mL in a volumetric flask

Attempt #1: Using strong organic acids to digest the blood cells and release the iron from the
hemoglobin in the blood

1. 2 mL of blood were added with 2 mL of water into a test tube

2. 5 mL of HNO3 was added to the test tube, and was placed into a boiling water bath
3. The sample (a pale yellow solution) was allowed to boil for about 10 minutes

4. 0.5 mL of conc. H2SO4 was added to the solution and boiled again

5. After about 5 minutes of boiling the tube was allowed to cool

6. Another mL of HNO3 is added, and then allowed to boil again

7. The test tube was again cool, under cold water from the tap, and 1 mL of distilled water was added
and 5 drops of hydrogen peroxide

8. This solution was then boiled again until white fumes appeared, and the test tube was then allowed
to cool

9. 5 mL’s of conc. HCl was added to the cool tube and then allowed to boil until a clear yellow solution
was obtained

10. The test tube was again cooled under the tap, and the solution was then transferred into a 125 mL
Erlenmeyer flask

11. 3 mL’s of water and 2 mL’s of 4% KI solution were added to the Erlenmeyer flask

12. The solution in the flask then had 1% starch indicator added and the solution was then titrated with
0.0125 M Thiosulfate solution

How this method failed: Since the procedure failed because instead of the solution turning the dark
bluish purple color when the KI and starch were added we spiked the blood with 2 mL of our iron
standard. When we performed the procedure with the spiked blood something crashed out of the
solution. We assumed this precipitate was the iron crashing out of solution so we filtered the
precipitate out of the solution. The filter paper was allowed to dry for 48 hours, and then massed. The
data we received is listed under Attempt #1. Also, while the Sodium Thiosulfate was not used
qualitatively, the standardization aspect of the experiment was included in the write-up because it was
used to qualitatively test for the presence of Iron in the filtrate (which will be elaborated upon in the
conclusion).
Data:

Standardizing Thiosulfate                                                      Molarity
                                                              Volume
                                                                               Na2S2O3
  Mass KIO3                                                Na2S2O3 per
                    1.0192                                                        per
     (g):                                                  Titration (L):
                                                                               Titration:
  Total moles                                                0.0145           0.1970738
                  0.004762617
    KIO3:                                                    0.01502         0.190251005
   Molarity                                                  0.0132          0.216482583
                  0.047626168
    KIO3:                                                    0.0145           0.1970738
  Moles KIO3                                                 0.01511         0.189117809
                  0.000476262
 per Titration:                                              0.01299         0.219982301
    Moles                                                                    0.20166355
 Na2S2O3 per      0.00285757                                                       M
  Titration:


                                                          Standard
                                                          Deviation:        0.013304


Attempt #1: Impromptu Gravimetric Analysis

 (Precipitate Crashing      Filter Paper+Precipitate   Filter Paper Alone              Precipitate Alone
          Out)              (after drying)
With Iron Standard          0.2197 g                   0.2164 g                        0.0033 g
With Spiked Blood           0.1434 g                   0.1364 g                        0.0070 g
Sample


This data still shows that the precipitate from the blood was not just iron. When the mass of the
precipitate of the iron standard is subtracted from the spiked blood sample the mass of iron in the blood
sample is still too high (0.0037 g or 3.7 mg) meaning something else crashed out of solution with the
iron, possibly proteins and other non-digested cell parts. The value of iron in blood should be 0.0946553
mg/mL (or 0.18927 mg per 2 mL sample). Our data turned out to be around 20 times larger, which leads
us to believe that there were contaminants in each trial.



Attempt #2: Using a reducing agent in the form of a mild acid and then the addition of a strong acid to
remove the iron from the hemoglobin of the cells, and then centrifuged to remove the proteins.

1. A solution of ascorbic acid was made by dissolving 0.8 grams of ascorbic acid in 50 mL of 0.1 M HCl

2. A solution of 30 wt% trichloroacetic acid was created by dissolving 3 grams of the trichloroacetic acid
into 10 mL of water
3. A 2 mL sample of blood was put into a test tube and 5 mL of the ascorbic acid was added to it, and
then agitated with a glass rod

4. The solution was then allowed to sit for 10 minutes

5. After the 10 minute period, 1 mL of the 30 wt% trichloroacetic acid was added to the solution

6. The test tube was then put into a centrifuge for approximately 8 minutes

7. The supernatant was then pipetted out of the test tube

8. The supernatant was then treated with 4% KI solution and the starch indicator

9. This was then titrated with the 0.0125 M Thiosulfate solution

Why this method went wrong: This procedure was actually created for a spectrophotometer reading of
the sample, but we believed we could have used this method to digest our cells a different way and
possibly produce results. We were going to use the ferrozine buffer in the suggested procedure to
obtain some data; however, we did not have the buffer in our prep room. We are not sure exactly why
this still did not provide a proper way to digest our cells and analyze our samples.

No data was garnered from this procedure.

Calculations

Standardizing Thiosulfate

Molarity Na2S2O3 = moles KIO3 X

Molarity Na2S2O3 = 4.762 E-4moles KIO3 X                                        = 0.1970738 M (stock
solution)



Checking the Masses of the Precipitate

Example for Trial using just Iron Standard

Mass Precipitate = mass (product + filter paper) – mass (filter paper)

Mass Precipitate = .2197 g - .2164 g = .0033 g = 3.3 mg
Conclusion

         The experiment did not go as planned. In Attempt #1, we had planned to oxidize all of the iron
in blood to its +3 state in order to perform an Iodometric titration with our Sodium Thiosulfate solution.
As mentioned in the procedure, this would have been accomplished by the addition of strong acids to
the blood samples, which would then have been boiled off of the analyte. However, the article gave no
specific amount of time to allow the boiling to occur; it simply said “until 2 mL of pale yellow solution
remained in the test tube.” However, due to time constraints, we were unable to allow it to boil all day
and still have time to run the titration. If given the opportunity to retry this experiment, we would allow
the samples to gently boil for a considerably longer time (≈4 hours), and we would take the time to
figure out what crashed out of the solution and for what reason.

        In Attempt #2, we tried to reduce the sample to its +2 state, and then remove all of the organic
materials via centrifugation. This method, however, was intended to be used for spectroscopic
analyses; we tried to use just the “preparation” steps (i.e. reducing the iron and separating it from the
organic materials), but stopped short of adding Ferrozine buffer. If we were able to conduct this
procedure again, we would try adding Ferrozine to obtain some sort of data, to see if it supports what
results we saw through our other methods.

         The only data we were able to collect was during Attempt #1 when a precipitate formed in our
solution before the addition of any Thiosulfate. This substance was then collected by vacuum filtration,
allowed to dry, and then massed. The filtrate was then titrated (qualitatively) with Thiosulfate to
determine if the precipitate was in fact iron. When the filtrate from the iron standard solution was
analyzed, the titration supported that there was in fact iron in the precipitate; however, the blood
sample did not act in this fashion, which leads us to believe that other material may have been present
in the precipitate and some of the iron remained in solution. It is also possible that not all of the iron
was removed from the cells before analysis.

				
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