# Colorimetric Fe Analysis

by Ulrich de la Camp and Oliver Seely [1] (Copied with kind permission and with no liability accepted for the current content.)

## Discussion

In the colorimetric analysis for Mn the concentration of the Mn is determined using the characteristic colour of the manganate (VII) ('permanganate') ion. However, few metal ions show such strong colours, particularly at low concentrations. Fortunately many highly coloured complexes can be formed from metal ions and organic or inorganic complexing agents. These complexes are the result of the interaction of a Lewis acid (the metal ion) and a Lewis base (the complexing agent). The ideal colour-forming reagent should be stable and selective (even specific) and react rapidly to form soluble, highly coloured complexes. The coloured complex should have a high absorptivity and be free from variations in colour due to minor changes in pH or temperature.

The application of colorimetric reagents is not a new technique but dates back nearly two thousand years. Around 60 A.D. Pliny the Elder in his "Natural History" recommended the use of nutgall as a reagent for the determination of iron in verdigris, which is a green pigment. Nutgall contains about 65-70% tannic acid which when combined with iron leads to the formation of a black iron/tannate complex.

In general organic colorimetric reagents are considerably more sensitive than are inorganic ones. They give more intense colours and are therefore frequently used for trace analyses. With many organic reagents, it is possible to determine concentrations at the ppm level. 2,2'-Bipyridyl (bipy), M = 156.20 g mol−1, forms an intensely red complex with iron(II) which may be exploited to determine iron concentrations in the ppm range. The reaction is:

3 bipy + Fe2+ ${\displaystyle \rightleftharpoons }$ Fe(bipy)32+

The figure below shows the structure of the reagent and the complex formed with Fe.

The molar absorptivity of the iron-bipyridyl complex is 8650 dm3 mol−1 cm−1 at the wavelength of maximum absorbance. The complex forms rapidly, is stable over the pH range 3 to 9, and may be used to determine iron(II) concentrations in the range of 0.5 to 8 ppm.

Iron(III), if present, must be reduced to iron(II) to produce the coloured species. A suitable reagent for this purpose is hydroxylamine hydrochloride, HONH3+Cl-. The reaction for this reduction is shown below:

2 HONH3+Cl- + 2 Fe3+ ${\displaystyle \rightleftharpoons }$ 4 H+ + 2 H2O + N2 + 2 Fe2+ + 2 Cl-

The concentration of iron in the sample could be calculated from Beer's Law however in this procedure we employ a different method. We will prepare a standard solution and compare absorbance readings of the sample and the standard solution. This technique minimises the effects of instrument and solution variation. Spectrophotometric methods are normally accurate to about ± 1%, i.e. to about three significant figures. Even though higher accuracy and precision can be obtained with more sophisticated instruments, in most cases an accuracy of ± 1%, at concentration levels of parts per million, is quite sufficient. The iron (II) ammonium sulphate standard that is used in the procedure is normally not considered a primary standard, however it is available in a purity greater than 99% and is therefore adequate for our purposes.

A major source of error in this experiment is misuse of the colorimeter. Before you take any measurements on this instrument read the instructions and commit them to memory.

## Experimental

### Preparation of the Original Fe Solution.

To an accuracy of ± 0.1 mg weigh out enough iron (II) ammonium sulphate, Fe(NH4)2(SO4)2 6 H2O, M = 392.14 g mol−1, to prepare 250 cm3 of a solution which is 0.00200 mol dm−3 with respect to that compound. Quantitatively transfer the salt into a 250 cm3 volumetric flask, add sufficient water to dissolve the salt, add 8 cm3 of 3 mol dm−3 H2SO4, dilute to the mark with distilled water and mix well. We shall call this the Stock Fe Solution. Pipette 10 cm3 of this solution into a 100 cm3 volumetric flask, add 4 cm3 of 3 mol dm−3 H2SO4 and dilute to the mark with distilled water and mix well. Label this solution as Original Fe Solution and calculate the concentration of Fe, in ppm, in this solution.

### Determination of the Absorbance of the Standard Fe Solution.

Before beginning this part of the procedure be sure to record the number of the colorimeter that you are using for this part of the analysis. The number of the colorimeter is found on a small blue tag on the front of the colorimeter. All further absorbance measurements must be made with the same colorimeter and the same cuvettes in order for this method to work. Discard both of the solutions in the 50 cm3 volumetric flasks. Thoroughly rinse both volumetric flasks and then prepare a new set of solutions from the Original Fe Solution using the same amounts according to the previous procedure. Label the solution containing the Fe as Standard Fe Solution and calculate its Fe concentration in ppm. Determine the absorbance of this solution at the wavelength of maximum absorbance previously determined. For a blank use the solution which does not contain Fe. Make at least three measurements. In each case reset the zero and the 100% transmission. Record both the percent transmission and absorbance values. Empty your cuvette and refill it with another portion of the same solution and again determine the absorbance value. Calculate the average of all six absorbance values.

### Analysis of city tap water.

If samples of city tap water are supplied in this experiment, you will determine the iron concentration in two samples. The concentration of iron in city tap water approaches the level of precision of this method because the concentration of iron in most water of southern California is very low. Iron pipes offer the most abundant source for the iron in our water. The Ksp of Fe(OH)3 is 4.0 x 10-38. Calculation yields a concentration of Fe3+ in neutral water to be so low as to be undetectable—on the order of one part iron per one quintillion parts of water (10-18). Still, whatever iron that finds its way into our water supply may end up in the form of a colloidal precipitate of ferric hydroxide. To bring that small amount of iron into solution we acidify tap water and boil it for two minutes, cool it to room temperature in ice and carry out the procedure now familiar to you.

Procedure

Calculations

The calculation of the Fe concentration of the unknown can be made by a comparison method. This, however, can only be done if the system adheres to Beer's Law in the range of concentrations involved. In the case of the iron-bipyridyl complex that range is 0.5 to 8 ppm. The appropriate relationship for the calculation of the Fe concentration in the Second Unknown Dilution is:

"A" is the absorbance, the subscript "S" refers to the absorbance and concentration, respectively, of the Standard Fe Solution while the subscript "SUD" refers to the absorbance and concentration of the Second Unknown Dilution. For the absorbance value of the unknown solution use the average of the three readings obtained for each sample taken. From the value obtained for [Fe]SUD, calculate the concentration of iron, in parts per million, of the original unknown Fe solution given to you by your instructor.

If city water was provided for this experiment, the equation above is used to determine [Fe] for the solution whose absorbance you measured. Since this solution was made to 50 mL but in the process you used 15 mL of reagents prepared with distilled water, [Fe] for the city water can be found through a simple modification of the equation above:

The Report

Report the following data:

1. Fe unknown number
2. Colorimeter number
3. Wavelength of maximum absorbance
4. Average absorbance of the Standard Fe Solution
5. Concentration of the Standard Fe Solution in ppm
6. Absorbances of the two samples of Second Unknown Dilution (SUD)
7. The average concentration of the Second Unknown Dilution in ppm.
8. Use the value given in 7, above, to calculate the original concentration in ppm Fe in the solution given you by your instructor. (To get to that point, you performed two dilutions: (a) 10:100 and (b) 10:50)
9. ppm iron in the two samples of city water, if provided for this experiment.
10. Pages in your lab notebook containing the pertinent data

Attach the original or a copy of the absorbance spectrum to the report sheet.

Questions on Colorimetric Iron

1. What is the name and structural formula of the ligand used in this procedure?
2. Draw the structural formula for the complex formed between Fe2+ and the ligand.
3. Why is hydroxylamine hydrochloride used in this procedure?
4. What is the formula of ferrous ammonium sulfate hexahydrate?
5. Why is sodium acetate used in this analysis?
6. Over what range of iron concentrations does the iron-bipyridyl complex obey the Beer-Lambert law?
7. Tea contains a significant amount of tannic acid. Given this fact explain why a cup of tea made with distilled water does not show the characteristic dark brown color of tea made with ordinary water.
8. List some characteristics of a good complexing agent for colorimetric analyses.
9. Calculate the ppm Fe in your Stock Fe Solution.