EXPERIMENT 4. Determination of Iron in Tap Water by Colorimetry

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1 EXPERIMENT Determination of Iron in Tap Water by Colorimetry Introduction Apart from zinc, iron is one of the most important transition metals in biochemistry. The most familiar biochemical compound containing iron is haemoglobin - the oxygen carrier in red blood cells. Iron is a necessary dietary element and is found in fruits, cereals, vegetables, and some meats. In certain cases (such as for pregnant women) it may be advisable to take an iron supplement (vitamin tablet). Iron is a common (and harmless) contaminant in ordinary tap water. In this experiment, the amount of iron in untreated natural water will be determined colorimetrically. The iron in solution must be present in the form of Fe 2+. To accomplish this, the iron tablet is first dissolved in acid. Then a reducing agent (hydroquinone) is added to ensure that all of the dissolved iron is in the form of Fe 2+. Hydroquinone Quinone The ph of the solution is adjusted by adding a buffer (sodium acetate, ph = 3.5). Finally, an intensely coloured species is formed by reacting the Fe 2+ with 1,10-phenanthroline solution (also known as the indicator, ferroin). This complex is very stable and the colour intensity does not change appreciably with time. Moreover, the colour is stable over a wide ph range (2 to 9). Phenanthroline has three fused rings containing two nitrogen atoms. Each nitrogen atom acts as a Lewis base by donating a pair of electrons to the Fe 2+, which acts as a Lewis acid. Fe (o-phen) [Fe(o-phen) 3 ] 2+ o-phenanthroline The intensity of the red coloration of the unknown is compared with that of a series of standards to estimate the amount of iron which is present in the unknown.

2 4-2 Materials: Reagents: ppm iron solution Unknown solution 1% hydroquinone in water 0.25% o-phenanthroline in water 2.5 % sodium acetate in water Buffers ph 4&7 for standardizing the ph meter Glassware: Apparatus: 2 cuvettes 7 volumetric flasks, 50 ml pipettes, 10 ml, 5ml, 2ml beakers for stock solutions spectrophotometers ph meter stirring bar Experimental Procedure Preparation of standard solutions Obtain your unknown sample of tap water from the stockroom. Using a 10 ml graduated pipette, transfer exactly 5.0 ml of the ppm iron standard solution into the 25 ml beaker. Test the ph of the solution with a calibrated ph meter (your demonstrator will show you how) and adjust its ph to 3.5 by dropwise addition of the sodium acetate solution using a medicine dropper. Count the number of drops (D) of sodium citrate solution needed. (No more than 50 drops should be needed.) Now pipette 5.0 ml of the standard ppm iron solution into a 50 ml volumetric flask. Using the same dropper as before, add the same number of drops of sodium citrate solution, followed by 1.0 ml of the hydroquinone solution and 1.5 ml of the o- phenanthroline solution. The resulting solution should be reddish in colour. Fill the volumetric flask to the line with deionized water mixing well. Allow the solution to stand for 10 minutes to permit the reaction to go to completion. Label this solution known 1. Prepare the remaining known solutions in the other 50 ml volumetric flasks according to the following table:

3 4-3 Solution Fe 2+ solution Acetate solution Hydroquinone (drops) known D known D x known D x known D x known D x known D x blank 0.0 D o-phen Each solution, made up to 50 ml with deionized water, should be allowed to stand for 10 minutes before use. Calculate the concentration of Fe 2+ in ppm in each solution for your report. Preparation of the unknown solution from tap water Pipette 5.0 ml of the unknown into a 25 ml beaker. Determine the number of drops of acetic acid solution needed to adjust the ph to 3.5 as described above. Take a new 5 ml aliquot in a 50 ml volumetric flask, add to it the same number of drops of acetic acid solution. Add to the volumetric flask 1.0 ml of the hydroquinone solution and 1.5 ml of the o-phenanthroline solution. Make the solution up to the mark with deionized water. Construction of a Calibration Curve and Analysis of an Unknown Record the spectra of all the above solutions on the Ultrospec 100 pro spectrometer over the range nm. Find λ max for each scan and determine the absorbance at that wavelength. Use a 1 cm plastic cell: 1. Turn on the Ultrospec 100 pro by pressing ON/OFF button. 2. Once the diagnostic t4ests have run and automatic calibration is complete, choose F2. This command will pull down the following menu: Repeat last operation Make a measurement Set up instrument 3. Press F2 to select Make a measurement. The Measurement option will open a new window with the following menu: Single/Multi λ

4 4-4 Cell density Select a method Scan 4. To measure Absorbance at a fixed wavelength, choose Single/Multi λ then Single λ. A new window will open: Abs/T% Set λ Scan Print 5. Press F1 to choose A or T%. 6. Set λ. λ limits are set by default. The instrument will scan between 330 and 830nm. 7. Press F3, Scan, to open the scanning window. 8. Start with the blank. Place the cuvette containing the blank solution in the cell holder and press the black button that indicates blank. The scan will take several seconds to complete. Because of the absence of iron in the blank, no peak should be observed. 9. Fill up the cuvette with the most concentrated sample and place it in the holder. Press the green button to scan sample. Determine the maximum absorbance (near 500 nm) using the arrows to move the cursor (left and right,up and down) on the peak. Record the value of the Absorbance at λ max. Repeat the measurements for each of the known solutions and unknown solution. 10. Construct a calibration curve by plotting the concentrations of the known iron solutions (x-axis) versus the corresponding absorbances (y-axis). Draw the best possible straight line through the data points. Use the absorbance of the unknown solution and the calibration curve to obtain the concentration of Fe 2+. The absorbance for the unknown solution should fall somewhere between known 1 and known 6. If this is not the case, consult your demonstrator.

5 Experiment 4 - Report 4-5 Name Date Results Working wavelength (λ max ) (nm) Cuvette path length (cm) (Only needed for question 2 below.) Calculate the concentration (molar) of Fe 2+ in each of your known solutions and enter them in the table below. solution Fe 2+ solution known known known known known known blank 0.0 Fe 2+ concentration absorbance at λ max Make a plot of concentration (x-axis) against absorbance (y-axis) for the above data and draw a best straight line through the points. Volume of tap water used Absorbance at λ max of the unknown solution Concentration of Fe 2+ read from graph (M) Calculation of the percentage of the concentration of iron in ppm in the tap water:

6 Questions In the experimental procedure, the number of drops of acetic acid needed to adjust the ph to 3.5 is determined on a separate aliquot which is subsequently discarded. Why? 2. What is Beer s Law? Use it, and your calibration data, to determine the molar extinction coefficient of the iron/o-phenanthroline complex at λ max.

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