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# Estimation of Iron (II) and Iron (III) concentration

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Introduction

The Estimation of Iron(II) and Iron(III) in a Mixture Containing Both Introduction - theory behind two methods and why I am choosing one of them There are several possible methods that could be employed to determine the concentration of Iron(II) and Iron(III) ions in solution: one of these is colorimetry. Colorimetry is the technique of using the depth of colour of a substance to measure its concentration. We use a colorimeter (see fig. 1 and 2) to measure the depth of colour. The machine is calibrated by checking a series of different known concentrations of solution; from the readings we construct a graph of absorbance (the percentage of light absorbed by the sample) against concentration - this is known as a calibration curve. We can then read off the graph the concentration value that goes with the absorbance value for our unknown solution. Figure 1: A colorimetry set up Figure 2: A colorimeter The machine may be set by the wavelength of the light involved. ...read more.

Middle

Another possible method would be a redox (reduction-oxidation) titration. Acidified potassium manganate(VII) is a popular standard solution for use in redox titrations and has the added benefit that it oxidises Fe2+ to Fe3+: MnO4-(aq) + 5Fe2+(aq) + 8H+(aq) --> Mn2+(aq) + 5Fe3+(aq) + 4H2O(l) The solution needs to be well acidified to provide the H+ ions shown in the equation. 2M sulphuric acid could be used to do this. The acid also inhibits the oxidation of Fe2+ to Fe3+ by the air. Without sufficient acid alternative reactions take place and the link between the MnO4- and Fe2+ will be lost. This titrant is self-indicating: a separate indicator is not needed. The manganate(VII) ion is bright violet, but it is reduced to the virtually colourless manganese(II) ion (it is actually very pale pink but looks colourless). When the iron(II) is finally used up, the manganate(VII) is no longer reduced, so the purple colour remains. ...read more.

Conclusion

Different metals would have different electrode potentials. More reactive metals ionise more easily and so they should leave more electrons on the metal before they reach equilibrium. In other words, the more negative the metal is, the more reactive it should be. If we could measure the potential difference between metal and solution we could use it to construct a table of reactivity from the different voltage readings. However, to measure this potential difference we would have to put a metal electrode into the solution and this would have its own electrode potential. A platinum electrode would be placed in a solution which is both 1 mol dm-3 in iron(II) ions and 1 mol dm-3 in iron(III) ions. Iron(II) has the lowest oxidation number and appears nearest the electrode: Fe3+(aq) + e- Fe2+(aq) Cyanide can be used to test if it is completely reduced. Calculations - justify each calculation Safety Table - remember concentrations Apparatus - put in table KMnO4 (0.0025M) Mixture containing unknown concentrations of Fe3+/Fe2+ ions in solution Ferrozine (0.2%) 1M Sulphuric Acid 0. ...read more.

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