This simplifies to:

This equation shows that for every five present, one is required for the reaction to be completed.

19.1ml of potassium manganate solution was used for titration. From this the amount of manganate ions used can be calculated as follows:

Amount of present =

Mass of present =

Since the volume of solution used was 25cm3, one-tenth of the total solution made from the five iron tablets, the above calculation shows that:

. Dividing this value by 5, we get 53.3mg of Fe present in each tablet.

We may also calculate the amount of Iron(II) Sulfate are present in each tablet:

Molar mass of Iron(II) Sulfate, FeSO4: 55.8+32+164 = 151.8g/mol

Amount of FeSO4 present:

/5 tablets

The mass of the five iron tablets was measured to be 1.552g±0.0001. Dividing this value by five gives 0.310g per tablet, or 310mg. This means that the percentage of Iron(II)Sulfate present in each tablet in mass is as follows:

Uncertainties calculation

Percentage uncertainty for pipette reading =

- Volume of iron (II) solution used = 25ml±0.05

Percentage uncertainty for burette reading =

- Volume of potassium manganate used = 19.1ml±0.05

Adding the percentage uncertainties together, the overall uncertainty is 0.462%. Converting this to the uncertainty of the total amount of iron sulfate present, we obtain:

Thus, the total amount of iron sulfate can be said to be 145mg±0.670

Literature value for mass of Iron Sulfate present in one tablet is 160mg. The percentage error of the experimental results can thus be calculated as follows:

Converting this into absolute error in the amount of iron, we obtain:

Thus, the total amount of iron sulfate with respect to its error value can be said to be 145mg±13.6.

Conclusion

In this investigation it was found through calculations that each of the iron tablet contained approximately 145mg of Iron (II) Sulfate. This is 15mg less than the listed amount of 160mg on the package, or 10% less than the literature value. This may have happened because some of the crushed tablets were left behind in the mortar. Another possible reason is that some of the iron in the iron (II) solution deposited in the bottom of the volumetric flask. This would mean that the pipette was filled with solution of less iron concentration than the average concentration. This would explain the smaller obtained value.

It was also found that the percentage error value was significantly bigger than the percentage uncertainty. This suggests that there was systematic error in this experiment. A possible way to solve this problem is to calibrate the burette and pipette next time before conducting the experiment.

Evaluation

This investigation has several flaws in its method. Potassium manganate (VII) solution was used in this experiment to titrate against the iron tablet solution. However, because it readily decomposes into manganese dioxide (MnO2), it is difficult to ever obtain a solution with exact concentration. The result is that the volume of manganate (VII) solution used might have been bigger than necessary due to decreased concentration as a result of decomposition. To minimize this from happening, the solution should be kept in a clean container and exposed to as little light as possible, as the solution decomposes faster under light and contaminated container.

The obtained result was smaller than literature value, and this may have been because some of the powdered tablet was left in the mortar. To improve the accuracy of this experiment, one should wash the mortar with sulfuric acid or distilled water, which can then be transferred to the volumetric flask. This would allow more of the iron tablet to be contained in the acidic solution, thus increasing the amount of iron present.

To extend this investigation, one can experiment on more types of iron tablets. Alternatively, a different content such as calcium or potassium can be investigated. This would require different types of tablets to be used.