Determination of Iron in Natural Water by Spectrophotometry Essay Sample
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Determination of Iron in Natural Water by Spectrophotometry Essay Sample
Aim: To determine the iron in natural water by spectrophotometry. Abstract: The iron in natural water was determined by utilizing spectrophotometric analysis. That was done by measuring the absorbance of five Fe(oPH)2+3 standards at 510 nm. From that information, a calibration curve was plotted and used to find the amount of Fe2+ that was in two unknown water samples based on the absorbance readings obtained with them at 510nm. The equation of the line was found to be y=0.1765x + 0.0705. It was then determined that there was no iron present in water sample A, while for water sample B, the iron was present in the proportions of 0.9037ppm, 1.614×10-5M and 9.037×10-3%. Introduction: Spectroscopy is the study of the interaction of light or electromagnetic radiation with matter. Spectrophotometry is any technique that uses light to measure chemical concentrations. Electromagnetic radiation is a form of energy when reacted with matter, can be absorbed, reflected or refracted, and how EMR reacts with matter depends on the properties of the material, based on the frequency, wavelength, absorbance etc.
The electromagnetic spectrum shows representative molecular processes that occur when light in each region is absorbed. The visible spectrum spans the wavelength range 380-780nm, so each region is absorbs at different wavelengths. The red-orange complex that forms between Iron (II) and 1,10-phenanthroline is useful for the determination of iron in water supplies. The reagent is a weak base that reacts to form phenanthrolinium ion in acidic media. The red-orange complex that forms between iron(II) and 1,10-phenanthroline (orthophenanthroline) is useful in determining iron in water supplies. The reagent is a weak base that reacts to form phenanthrolinium ion, phenH+, in acidic media. A commonly used method for the determination of trace amounts of iron involves the complexation of Fe2+ with 1,10-phenanthroline (phen) to produce an intensely red orange colored complex: Fe2+ + 3phen Fe(phen)32+. Since the iron present in the water predominantly exists as Fe3+, it is necessary to first reduce Fe3+ to Fe2+.
This is accomplished by the addition of the reducing agent hydroxylamine. An excess of reducing agent is needed to maintain iron in the +2 state (because dissolved oxygen will reoxidize Fe2+ to Fe3+). Fe2+ is quantitatively complexed by 1,10-phenanthroline in the pH range from 3 to 9. Sodium acetate is used as a buffer to maintain a constant pH at 3.5. If the pH is too high, the Fe2+ will be oxidized to Fe3+; if the pH is too low, H+ will compete with Fe2+ for the basic 1,10-phenanthroline (to form phenH+). Either way, complete complexation won’t be achieved. The determination of the iron-phen complex is performed with a spectrophotometer at a fixed wavelength of 510nm using external calibration based on iron standard solutions. In the spectrophotometer, light is passed through a monochromator (a prism, a grating, or even a filter) to select a range of wavelength and some of the light may be absorbed by the sample therefore giving the transmittance, which is the fraction of the original light that passes through the sample and has the range 0 to 1. Absorbance, sometimes called optical density, is the heart of spectrophotometry as applied to analytical chemistry Beer-Lambert’s law, A=ebc , where the concentration of the sample, M, path length , cm, quantity (epsilon) is called the molar absorptivity.
Molar absorptivity is the characteristic of a substance that tells how much light is absorbed at a particular wavelength. Both a and e depend on the wavelength of electromagnetic radiation. Attenuation of radiation as it passes through the sample leads to a transmittance of less than 1. Besides absorption by the analyte, several additional phenomena contribute to the net attenuation of radiation, including reflection and absorption by the sample container, absorption by components of the sample matrix other than the analyte, and the scattering of radiation. To compensate for this loss of the electromagnetic radiation’s power, we use a method blank.
Method: Standard iron solutions, sodium acetate solution, 10% hydroxylamine hydrochloride, 1-10 phenanthroline and solvent blank solutions were prepared before hand to use throughout the experiment. 2.00, 4.00, 6.00, 8.00 and 10.00 mL of iron stock solution were pipetted into five 100mL volumetric flasks. To each flask, 1mL of 10% hydroxylamine hydrocholride, 10 mL of sodium acetate and 10mL of 1,10-phenanthroline solution was added. The mixture was allowed to stand for 10 minutes then made up to the mark with distilled water. The absorbance of all five standards solutions were determined with respect to the blank at 510 nm. 10.0mL of water sample was transferred to a 100mL volumetric flask, and treated exactly the same way as the standards, measuring the absorbance with respect to the blank. Results:
Concentration of Fe(oPH)2+| Absorbance at 510nm|
Unknown| Absorbance at 510nm|
Table 1 showing the absorbance obtained at a wavelength of 510nm, using a OHAUS spectrophotometer, for concentrations of Fe(oPH)2+ of (1,2,3,4,5)ppms and two unknown water samples A and B. Calculations:
(i) Equation of the line: y=0.1765x + 0.0705. Sample A absorbance -0.007, hence, that value does not have to be substituted in the equation as it is a known fact that a negative absorbance reading means there is no iron analyte present.
(ii) Since a (1/10) dilution was undertaken, the absorbance reading for sample B could also be expressed as (0.023 x 10) = 0.23. Substituted in the above equation we get: 0.23= 0.1765x + 0.0705
0.23- 0.0705= 0.1765x
x=0.1765/0.1595= 0.9037ppm, or 0.9037mg/L
(III) Ar of Fe = 56g/mol
moles of Fe2+= mass/molar mass= 0.9037 x 10-3g/56g/mol= 1.614×10-5M Fe2+
(iv) density of water= 1g/mL or 1g/cm3
mass of water= density x volume= 1 x 10= 10g
(v) %Fe2+ in water sample B= = 0.9037 x 10-3g Fe2+ x 100/ 10g water= 9.037 x 10-3 % Fe2+
Discussion: The concentration of standard A was 1 ppm and its absorbance measured at 510 nm was 0251. The concentration of standard B was 2 ppm and its absorbance measured at 510nm was 0.424. The concentration of standard C was 3 ppm and its absorbance measured at 510nm was 0.587. The concentration of standard D was 4 ppm and its absorbance measured at 510nm was 0.785. While the concentration of standard E was 5 ppm and its absorbance measured at 510nm was 0.482. The trend observed with the concentration of the first four Fe(oPH)32+ standards and the absorbance readings obtained with them was that as the concentration increased, so too did the absorbance readings measured. That meant that concentration was directly proportional to absorbance (Aᾳ C). Hence, Beer-Lambert’s law was observed. The concentration values as well as the absorbance readings at 510nm for the Fe(oPH)32+ standards were then used to plot a graph of absorbance versus concentration, with the dependent on the x- axis.
The graph plotted yielded a straight line and the equation of the line was found to be y=0.1765x + 0.0705. It is important to note at this point that 510 nm was the used wavelength because it represented λmax for the Fe(oPH)32+ complex. In other words, it is the wavelength at which the complex absorbs best and therefore the peak of its absorbance spectrum would be located at that wavelength. Subsequently the equation of the line mentioned above was utilized by using the absorbance readings for water samples A and B, which were -0.007 and 0.023 respectively. Since the absorbance readings obtained for sample A was negative, that meant that there was no iron present for sample A. While for sample B, the value was positive and since A (1/10) dilution was used calculations was done, using and absorbance value of (0.023 * 10= 0.23).
It followed that the iron was present in water sample B in the proportions of 0.9037ppm, 1.614×10-5M and 9.037×10-3%. Hence, iron was present in trace amounts in sample C. This experiment was not short of errors though, as there was one questionable results that was encountered in this experiment. For the absorbance readings obtained for the 5 ppm standards was 0.482, which was contradictory to Beer’s law. That may have been as a result of the formation of iron salts such as phosphates. As the acid buffer may have failed to keep the pH at the optimum of 3.5. It may have been also as a result of using wrong proportion of reagents by experimenters when making up that standard. Hence, when conducting this experiment, suitable reagents must be used in order to maintain optimum reaction vessel conditions. The experimenters must also ensure that he/she follows all the instructions given with regards to making up standard to prevent undesirable results.
Conclusion: Spectrophotometric analysis was successfully utilized to generate a calibration curve for Fe(oPH)32+ standards. The equation of the line was found to be y=0.1765x + 0.0705. In addition, that equation was manipulated to determine the amount of Fe2+ in the two water samples. It was therefore determined that waste sample A had no iron present in it, but sample B contained 0.9037ppm, 1.614×10-5M and 9.037×10-3% Fe2+.
Skoog and West, Fundamentals of Analytical Chemistry, 2nd Ed., Chapter 29. Vogel, A Textbook of Quantitative Inorganic Analysis, 3rd Ed., p. 294, 310 and 787.