Introduction

Oxidation numbers are "bookkeeping symbols". Each atom in an equation can be assigned an oxidation number according to certain rules. If the oxidation number of an atom increases as you go from the left side to the right side if an equation, oxidation has occurred (electrons have been lost); if the oxidation number decreases, reduction has occurred (electrons have been gained).

(1) MnO₄-+ 8H⁺ + 5e- Mn⁺² + 4H₂O

In the above reduction half-reaction, manganese has undergone a decrease in oxidation number from +7 to +2. Whereas, in the following oxidation half-reaction, each Iron atom has undergone an increase in oxidation number. From +2 to +3.

(2) Fe ⁺²  Fe ⁺³ + 1 e-

Oxidation must occur along with reduction. These reactions are called redox (reduction/oxidation ) reactions. The number of electrons lost and gained in the half-reactions must be equal. The overall redox reaction becomes:

(3) MnO₄+ 5 Fe⁺² + 8H+ Mn⁺² + 5 Fe⁺³ + 4 H₂O

In Part I of the experiment you will standardize a KMnO₄ solution by titrating it against a Ferrous Ammonium Sulfate (Fe(NH₄)₂(SO₄)₂·6H₂O) solution where the number of moles of Ferrous Ammonium Sulfate is precisely known. Using a burette you will slowly add the KMnO₄ solution to the Ferrous Ammonium Sulfate solution. Reaction (3) will occur. KMnO₄ acts as its own indicator. That is, at the start of the titration, the deep violet color of MnO₄-will be lost because it is changing to Mn⁺², but as soon as there is no more Fe²⁺ in the solution for the MnO₄-to react with, the MnO₄-will remain in the reaction solution. The pink color of the dilute MnO₄- solution indicates the end of the reaction. This is known as the "end point" or equivalence point of the titration.

Knowing the mass of Ferrous Ammonium Sulfate(s) to the nearest 0.0001g, you can precisely determine the number of moles of Ferrous Ammonium Sulfate(s). This will be equal to the number of moles of Fe²⁺ ions in solution. Using equation (3), one can then determine the number of moles of MnO₄-that have reacted during the titration. Knowing the milliliters of MnO₄-solution used in the titration, the molarity of the MnO₄-can be calculated as molarity is equal to the moles of solute (MnO₄-) per liter of solution. This will be equal to the molarity of the KMnO₄ solution as one mole of KMnO₄ dissolves in water to give one mole of MnO₄-. You now know the precise molarity of the KMnO₄ solution. That is another way of saying that you have standardized the KMnO₄ solution.
Part II of the experiment is very similar. We will us the standardized KMnO₄ solution to determine the % by mass of Fe in a sample. The concentration of the KMnO₄, remains the same as Part I, but the amount of Fe⁺² present in the sample is unknown. The titration is performed in the same way.

(3) MnO₄- + 5 Fe⁺² + 8H⁺ Mn⁺² + 5 Fe⁺³ + 4 H₂O

Since we know the precise molarity of the MnO₄- solution from Part I.  Multiplying liters of MnO₄- solution used in a titration times moles of MnO₄- per liter of solution (its molarity) we obtain moles of MnO₄- that have reacted. Using equation (3) we can then determine moles of Fe⁺². This will allow us to calculate the number of grams of Fe in the sample. Finally, knowing the mass of the sample, we can calculate the percent iron in the sample used in the titration, which is the whole point of the experiment.

Equipment  

3 · 250 mL Erlenmeyer flasks	Buret
buret clamps	ring stand
analytical balance	100 mL graduated cylinder

Chemicals
student Fe(NH₄)₂(SO₄)₂·6H₂O, 0.02 M KMnO₄, 3 M H₂SO₄, unknown Fe⁺² sample, concentrated H₃PO₄.

Spill/Disposal
Fe(NH₄)₂(SO₄)₂·6H₂O, unknown Fe⁺² sample, all reacted chemicals and small 0.02M KMnO₄ spills are Spill/Disposal A.
3 M H₂SO₄ & concentrated H₃PO₄ are Spill/Disposal B1.
Any 0.02M KMnO₄ solution left in the buret is to be returned to the stock bottle.

Procedure

Part I

1. Rinse a buret thoroughly with tap water followed by several rinsings with distilled water. If the buret is clean, no water droplets will remain inside the buret. Rinse the buret carefully once with 5mL portions of 0.02 M KMnO₄. Now, fill the buret with the 0.02 M KMnO₄ solution. Note that its concentration is given only one significant figure. It must be standardized to get a more precise concentration. Be sure that the tip is filled.

2. Weigh out two 0.5g samples of dried Ferrous Ammonium Sulfate Fe(NH₄)₂(SO₄)₂·6H₂O on an analytical balance to the nearest 0.0001 g. Record the weights on the data sheet. Place each sample in a 250-mL Erlenmeyer flask. Be sure to get, the entire sample, into the flask. Dissolve each sample in 75mLs of distilled water. Add 5mL of 3 M H₂SO₄, and 3mLs concentrated H₃PO₄ (85%) to each flask. (The Fe⁺³ that will be produced during the titration forms a colorless complex with the PO₄- ³ ions. This simplifies the detection of the end-point.)

3. Record the initial reading on the buret (read the bottom of the meniscus at eye-level) to the nearest 0.02mL. Start to add the KMnO₄ (aq) solution. When the violet color of the MnO₄-ion in the reaction does not disappear quickly, add the solution slowly. Towards the end of the titration, the solution should be added one drop at a time. When a faint pink color persists for 30 seconds, with constant swirling, the end-point has been reached. A white piece of paper under the Erlenmeyer flask will aid in detecting color changes.

4. Record the level of KMnO₄ solution in the buret to 0.02mL.

5. Repeat steps 3 and 4 with the second flask containing dried Ferrous Ammonium Sulfate. Begin with the buret filled nearly to the top of the graduated part.

6. Calculate the moles of Ferrous Ammonium Sulfate and the molarity of the KMnO₄ solution using the data from each titration. Your molarities for KMnO₄ should agree within 1%. If they don't, do a third titration.

Part II

1. Weigh out two 0.4g samples of the unknown mixture containing Fe⁺² ion an analytical balance to the nearest 0.0001 g. (Record unknown number) Place each sample in a 250mL Erlenmeyer flask. Dissolve each sample in 75mLs of distilled water. Add 5mL of 3 M H₂SO₄, and 3mLs concentrated H₃PO₄ (85%) to each flask.

2. Titrate the unknown solution with the now standardized KMnO₄ solution. The procedure is the same as in the above titrations. Again, towards the end of the titration, the 0.2M KMnO₄ has to added one drop at a time.

3. Calculate the moles of Fe⁺² present in the unknown sample and the percent by weight of iron in the sample. Again your results should agree within 1%.

4. Empty unused 0.2 M KMnO₄ back into the reagent bottle. Rinse your buret several times with water.

Disposal All contents of the reaction flask may be disposed of into the sink. Any unused KMnO₄ solution (in the buret) must be emptied into the 0.2 M KMnO₄ reagent bottle.