Determination of Iron in Water ENVS4010

Theory 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) http://www.meta-synthesis.com/webbook/15_redox/redox.html LHS: Fe2+ RHS: Fe3+

Chemistry Reduction of iron: 2Fe3+(aq) + 2NH2OH(aq) + 2OH-(aq) → 2Fe2+(aq) + N2(g) + 4H2O(l) Complexation by 1,10-phen Fe2+(aq) + 3(C12H2N2)3(aq) → [(C12H2N2)3Fe]2+(ag)

[(C12H2N2)3Fe]2+

Questions What is a chromophore? Why is it necessary to wait 10 minutes before reading the spectrophotometer? 3. What is a chelating agent? A chemical group capable of selective light absorption resulting in the coloration of certain organic compounds. To allow time for the formation of he complex (a) The standard recipe used in this exercise uses 3.0 mL of 0.25% o- phenanthroline solution in 100mL of final solution to generate the red iron- phenanthroline complex which is determined colorimetrically. What is the initial molarity of o-phenanthroline in the diluted solution? [0.25% means 2.50 g. of o- phenanthroline (molar mass 180) per liter of solution.] (b) What is the maximum concentration of iron-phenanthroline complex that can be formed from this amount of o-phenanthroline, assuming complete complexation? (c) What does this concentration of iron correspond to in mg/L? (d) Is this enough for the amounts of iron (up to 20mL of stock solution) we are using? (i.e.: What is the limiting reagent in these reactions?) A chemical compound in the form of a heterocyclic ring, containing a metal ion attached by coordinate bonds to at least two nonmetal ions.

Method 1. Preparation of Standard Solutions Prepare four iron (II) standards having the following concentrations: 0.5, 1.0, 2.0, and 5.0 ppm. Pipet 0.5, 1.0, 2.0, and 5.0 mL of 100 ppm stock solution into 100 mL volumetric flasks and dilute to the mark with DI water. 2. Obtaining a Beer’s Law Plot Transfer a 5 mL aliquot of the 0.5 ppm iron standard to a 125 mL Erlenmeyer 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 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. Use the same number of drops of sulfuric acid and sodium citrate for the remaining three standards and a reagent blank, followed by 3 mL of 0.3 % 1,10-phenanthroline and 1 mL of 10 % hydroxylamine. Mix well. Using water as the reference, measure the absorbance of each standard at 512 nm. 3. Analysis of samples Natural and tap water samples often have less than 0.5 ppm iron. Set the UV/VIS spectrometer to read at 512 nm. Determine the iron in several environmental water samples. Use 5 mL samples and treat them the same way as the standards, adding sulfuric acid initially (if necessary), followed by citrate buffer, reducing agent and the indicator. NOTE: All iron solutions should be discarded into a “Heavy Metal” waste container. Dissolve 0.7022 grams of ferrous ammonium sulfate, hexahydrate in distilled water. Dilute to 1.00 L. This solution is 100 mg/L Fe+2 (same as 100 ppm).

Calculations 1. For the calibration curve, select the x-y scatter chart that shows only the data points, then perform a best-fit line of your data without forcing the line through the origin. Determine the slope, y-intercept, and the correlation coefficient; comment on the linearity of the plot. 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. 3. If duplicate determinations were done, report average values and the individual values of concentration and discuss the reproducibility. 4. For the instrument you used, what would you estimate to be the lowest concentration of iron that would be reliable?

A = εbc

Questions 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 10 minutes on step 6? 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 acetate. 5. The analytical sensitivity of the method depends on the slope of the calibration curve. Was your method sensitive for the determination of iron in these samples? Why? 1. To minimize effects of interfering species. 4. 1,10-Phen complexes Fe2+, Hydroxylamine hydrochloride reduces ferrous to ferric ions. Sodium acetate buffers to pH 3.5

Field Instrumentation Palintest - Ferrometer 1000 1,10 phenanthroline test 0.02 to 5.00 mg/L Fe http://www.palintestusa.com

References 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.