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

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. Visible light wavelengths vary from around 400 nm (blue) to 700 nm (red).

Colorimetry is relevant to our investigation as Fe(II) will go from pale green (almost colourless) to form a violet complex ion with the bidentate ligand ferrozine (see fig.3). The formula of this complex ion is [Fe(ferrozine)3]2+. A solution containing ferrozine would appear coloured to us because white light enters and violet light is not absorbed (otherwise there would be no violet to see). It is the other colours that are absorbed. So if we were running some samples of ferrozine in a colorimeter we might put in a green filter (as green is violet’s complementary colour) and measure how much ...

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Figure 3: Ferrozine

The Fe3+ will turn red when Thicyanate is added.

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. As the titration should have been done to the nearest drop, there should be no more than a drop of purple manganate(VII) in the titration mixture. Consequently, it does not look purple, but appears pink. The end-point is the first permanent pink colour (see fig. 4).

Figure 4: the subtle colour change as Fe(II) is used up and manganate(VII) ceases to be reduced

→

The Iron(III) concentration can also be measured using this titration. We first reduce the iron(III) solution to iron(II). This can be done using a mixture of excess zinc metal and acid. Bubbles of hydrogen will be seen given off as the zinc reacts with the acid. After reduction is complete the remaining zinc is removed by filtration. The iron(II) concentration is now measured by titration with potassium manganate(VII) solution.

Using electrode potentials would also tell us if the reaction is feasible: the difference in charge between the metal and its solution means that there is a potential difference set up between them, which would be measured in volts. This potential difference is called the electrode potential of the metal. 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

Method – justify each step, third person, informative, point by point

Treatment of Results - questions

Bibliography