Titrations of Permanganate

       Titration is a common method for determining the amount or concentration of an unknown substance. The method is easy to use if the quantitative relationship between two reacting substances is known. The method is particularly well-suited to acid-base and oxidation-reduction reactions. In this experiment, you will conduct two separate redox titrations using a standardized permanganate solution. In the first titration, you will be trying to find the % hydrogen peroxide in a commercially sold solution. In the second titration, you will be trying to find the % iron in an unknown iron salt.
        Permanganate ion is a powerful oxidizing agent, especially in acidic solution, which can be used to analyze (by titration) solutions containing many different species. In these titration reactions, the intensely colored MnO₄^- ion is reduced to form the colorless Mn2+ ion.
An advantage of using the permanganate ion in the titration of colorless unknown solutions is that it is "self-indicating". As long as the reducing agent remains present in the sample, the color of MnO₄^- quickly disappears as it is reduced to Mn²⁺. However, at the endpoint, all the reducing agent has been used up so the next drop of MnO₄^- solution is sufficient to cause an easily detected color change, colorless ( faint, permanent peach/pink. So we know that at the endpoint, the oxidizing agent (MnO₄^-) and reducing agent (H₂O₂ or Fe²⁺) have reacted in exactly in proportion to their stoichiometry in the balanced redox equation. If we know how much of the oxidizing agent we added, then we can figure out exactly how much reducing agent was present in the unknown.

If you know how much of the reducing agent (H₂O₂ or Fe²⁺) was in the solution, you can calculate the mass of the reducing agent present, then calculate the mass% of species in the orginal sample ([mass Fe/mass of sample] for example)

Procedure:
0.1 M MnO₄^- solution
      Prepare 250.0 mL of a "standard" 0.1xxx M solution of MnO₄^- using KMnO₄(s) as the permanganate source. You should calculate the amount of KMnO₄ you will need before coming to lab. You will need to know the exact mass that got into the flask, but the 250 mL flask is too heavy to put on the analytical scale. When diluting the solution in the flask, add about 125 mL of water and 30 mL of 3 M H₂SO₄, dissolve the solid completely, and then dilute to volume. Once made, put your solution in DRY water bottle you brought to class. Cap tightly and label with your name and concentration.

Titration of an Unknown Hydrogen Peroxide Solution
       In a 125 mL place a known MASS (~5 g) of commercial hydrogen peroxide. To it add ~50 ml of DI water and 20 mL of 3 M H₂SO₄. Fill your buret (after 2x 5mL rinse, remember) with your MnO₄^- solution. Now slowly begin titrating the H₂O₂ solution while it is continously being stirred by gently swirling the flask. Continue titrating until you see the color of MnO₄^-begin to persist locally in the solution, at which point, you should slow down to dropwise additions. Continue until one added drop of MnO₄^- solution produces a faint peach/pink color that lasts at least 30 seconds. This is the first excess MnO₄^- which is not being reduced by the H₂O₂. Determine how much MnO₄^- was added in the titration and then calculate the mass% of H₂O₂ in your sample. Repeat runs until you have 3 that are within 2% of each other. You will report average and standard deviation from 3 ‘good’ runs.

Titration of an Unknown Iron Salt
1. Obtain approx. 1-2 g of the unknown iron salt, add ~50 ml of DI water and 15 mL of 3 M H₂SO₄. Titrate as before, getting at least three ‘good’ runs. The unknown salt could be any one of the following: Iron (II) sulfate, iron (II) nitrate, iron (II) chloride, or iron (II) sulfate heptahydrate (7 waters for every 1 iron (II) sulfate). You will determine the mass % of Fe in your solid sample and it should match of of these choices.
2. When finished, clean and rinse both the volumetric flask and buret thoroughly because MnO₄^- solutions will cause bad glass stains. Rinse with an acidic solution of H₂O₂ which will be provided by your instructor. Place about 50 mL of the solution in a flask. Swirl the solution until the flask is clean and then transfer it to another piece of stained glassware. When you are done using it, pass it on to another group as it can be ‘recycled’

Lab report: Filled out data sheet (see class blog) complete with sample calculations sheet (seperate) and commentary.

Prelab Questions: (worth 8 points)

1. What mass of KMnO₄ will you need to make 250 mL of a 0.1 M MnO₄^- solution?
2. In this lab, you will be reducing MnO₄^- to Mn²⁺ in the presence of H⁺. The reduction half reaction is
MnO₄^-(aq) + 8H⁺(aq) + 5e- → Mn²⁺(aq) + 4 H₂O(l)
The H₂O₂ is being oxidized into H₂O and O₂. That half reaction is
H₂O₂(aq) → O₂(g) + 2 H⁺(aq) + 2 e-
Write the overall (balanced) REDOX reaction that is occurring in H₂O₂ titration portion of this experiment.
3. In the other section of this experiment, you will be oxidizing Fe²⁺ into Fe³⁺. The oxidation half reaction for that should be easy to figure out. Write the overall (balanced) REDOX reaction that is occurring in the iron salt portion of this experiment.
4. 16.4 mL of a 0.104 M solution of KMnO₄ were needed to titratite (to endpoint) a 5.013 g sample of H₂O₂ solution. What is teh mass % of H₂O₂ in the sample?

--- You will need to bring in an EMPTY water bottle for storing your KMnO₄ solution that you will make. This should be a 'disposable' bottle as you will not want to use it after lab.---

Last updated by MEO :25Feb08