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Determination of Iron in Water by 7326kbw


									Determination of Iron in Water


•   Fe2+ (ferrous) complexed with 1,10 phenanthroline
•   Orange color
•   Spectrophotometer measures light absorbed
•   Selective for Fe2+
                                           Part A

•   The test for the thiocyanate ion is to
    add iron (III) nitrate to a solution. If the
    solution turns blood-red upon the
    addition of iron(III) nitrate this
    indicates the presence of the
    thiocyanate ion

•   Fe3+(aq) + 6 SCN-(aq) → [Fe(SCN)6]3-(aq)

                                                    LHS: Fe2+   RHS: Fe3+
1.     Reduction of iron:

     2Fe3+(aq) + 2NH2OH(aq) + 2OH-(aq) → 2Fe2+(aq) + N2(g) + 4H2O(l)

2.     Complexation by 1,10-phen

            Fe2+(aq) + 3(C12H2N2)3(aq) → [(C12H2N2)3Fe]2+(ag)

            Fe       N



1.   What is a chromophore?
     A chemical group capable of selective light absorption
     resulting in the coloration of certain organic compounds.

2.   Why is it necessary to wait 10 minutes before
     reading the spectrophotometer?
     To allow time for the formation of he complex

3.   What is a chelating agent?
     A chemical compound in the form of a heterocyclic ring,
     containing a metal ion attached by coordinate bonds to at least
     two nonmetal ions.
           Making the Standard Fe Solution

•   To determine ppm iron, you need to know that ferrous ammonium sulfate is
    Fe(NH4)2(SO4)2 ·6H2O and has a molecular weight of 392.14 g/mol. Note
    that ferrous ammonium sulfate is sometimes abbreviated as “FAS”

•   First calculate the % Fe in FAS (= 55.84/392.14 = 14.24 %)

•   Then divide the desired conc. by the % Fe in this compound:

•   100 ppm standard required:

    100 mg Fe per liter = 702 mg FAS / L

    = 351 mg FAS in 500 mL
NOTE: All iron solutions should be discarded into a “Heavy Metal” waste container.
1. Obtain your unknown.

2.   Standard 100 ppm Fe Solution: Obtain 100 mL of 100 ppm stock ferrous ammonium sulfate (Fe(NH4)2(SO4)2
     ·6H2O) solution (351 mg in 500 mL) in a clean conical flask. Be certain to note the mass of ferrous ammonium
     sulfate used and the final volume of the solution.

3.   Preparation of stds: Prepare five standard solutions of iron (II) having the following concentrations 0 (reagent
     blank), 0.5, 1.0, 2.0, and 5.0 ppm. Pipette 0.5, 1, 2 and 5 mL aliquots of the standard iron solution into 100 mL
     volumetric flasks and dilute to the mark with DI water.

4.   Obtaining a Beer’s Law Plot: Transfer a 5 mL aliquot of the 0.5 ppm iron standard to a 125 mL flask and test
     the pH with test paper. If greater than 4.5, add enough 0.2 M sulfuric acid dropwise to bring the pH to about 3.5,
     counting the number of drops. Again counting drops, add sodium citrate (259 g/L) buffer to bring the pH to about
     4.5. Pipet 1 mL of 10% hydroxylamine hydrochloride and 3 mL of 1,10-phenanthroline into the sample, mix, and
     allow 5 minutes for color development. DO NOT PUT YOUR PIPETTES INTO THESE SOLUTIONS; pour a small
     amount into your beaker and pipette from this. Be sure to add the reagents in the order shown here.

5.   Use the same number of drops of sulfuric acid and sodium citrate for the remaing standards, followed by 1 mL of
     10% hydroxylamine hydrochloride and 3 mL of 1,10-phenanthroline.

6.   Adjust the spectrometer to 512 nm, allow 30 mins for the instrument to warm up. Using DI water as the reference
     blank to zero the instrument, measure the absorbance of each standard (including the reagent blank)

7.   Analysis of Samples: Natural and tap water samples often have less than 0.5 ppm iron. Determine the iron in
     several environmental samples.

     Treat samples the same way as the standards, adding sulfuric acid initially, if necessary, followed by citrate buffer,
     reducing agent and indicator. Use 5 mL samples and adjust the pH for each sample individually.
1. Draw a calibration curve, then perform a best-fit line of your data without forcing the line through
    the origin.

Outline of Calculations:

1.   Describe how the 100 ppm Fe standard solution was made.

2.   You need to do dilution calculations for each individual iron standard solution, seeking to
     determine the concentration of the final solution (as tested in the spectrometer) in ppm Fe
     (mg Fe /L). The basic equation is C1 x V1 = C2 x V2

3.   Use the calculated concentrations of the diluted standards as the x values for graphing your
     calibration curve, paired with the absorbance values (as the y value) for those solutions.

4.   Use the equation of your best fit line on the calibration curve to determine the concentration of iron
     (in ppm Fe) in your unknown sample. Use correct sig. figs. That value (the concentration of iron in
     the diluted solution, as measured in the spectrometer is the final value you are expected to report.

5.   If duplicate determinations were done, report both individual and average values.

• A = εbc                0.5





                               0.5   1.0   1.5   2.0        2.5        3.0   3.5   4.0   4.5


1. Why use a “reagent blank” and not just distilled water to zero the spectrometer?

2. What would the effect be of waiting 30 minutes instead of 5 minutes for step 4?

3. Today’s unknown samples are simple iron dissolved in water. If our sample were
more complex, it might contain other compounds, some which might be colored,
including some which might absorb light at the same wavelength as the iron
phenanthroline complex. What problem would this cause? How could we modify our
procedure to correct for this problem?

4. State the function of each reagent in this experiment: 1,10-phenanthroline,
hydroxylamine hydrochloride, sodium citrate.

5. For the instrument you used, what would you estimate to be the lowest
concentration of iron that would be reliable?
           Field Instrumentation

• Palintest - Ferrometer 1000
• 1,10 phenanthroline test
• 0.02 to 5.00 mg/L Fe

• E. B. Sandell, Colorimetric Determination of Traces
  of Metals, 3rd Ed. Interscience Publishers, Inc., New
  York, 1959.
• D. A. Skoog, D. M. West, F. J. Holler, and S. R.
  Crouch, Analytical Chemistry: An Introduction, 7th
  ed., Chapters 21 and 22, pp. 547-592.

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