The reduction of an oxidizing agent often produces different products, depending on the pH of the solution in which the reaction is carried out. You will find out (by titration) in this experiment what happens when permanganate ion is reduced in acid, neutral, and basic solutions. Permanganate is a strong oxidizing agent and as such it picks up electrons. Because the product formed in acid, neutral, or basic solutions may be different, the number of electrons picked up per MnO4- may differ. You can determine the amount of reducing agent used and the number of moles of MnO4- reduced, and thus you can decide the number of electrons each MnO4- has picked up. Since manganese has an oxidization state of 7+, you can then determine the oxidation state of the manganese in the product in each solution.The reducing agent you will start with is sodium bisulfite, NaHSO3. The sulfur will change its oxidation state from 4+ to 6+ in all three cases, however the sulfur-containing species are different for each case.
Acid solution: the bisulfite ion becomes H2SO3 which then is oxidized to HSO4-.
Neutral solution: the bisulfite ion is oxidized to SO42-
Basic solution: the bisulfite ion becomes SO32- and is oxidized to SO42-
Procedure:
Acid:
Set up the buret after rinsing once with distilled water and twice with 5-mL portions of the permanganage solution. Then fill the buret with the standard solution of KMnO4. Be sure you record the molarity.Measure out 25.00 mL of NaHSO3 solution, also recording its molarity, and put it into a clean 150 mL beaker. Add about 5 mL of dilute sulfuric acid. Titrate with stirring, to the appearance of pink color that persists for at least 30 seconds. The pink color, due to excess MnO4-, can be seen better if you put a white piece of paper under the beaker.
Neutral:
Refill the buret, and repeat the titration exactly as in the acid solution, except with the omission of the acid. During the titration there will be formation of a brown precipitate which may first appear yellow in color. This is not the end point. Add KMnO4 until a permanent pink color persists. To facilitate seeing the pink color, you can use a capillary tube to periodically remove a small portion of the solution. Look at the solution in the capillary. You will find it relatively free of precipitate. Blow the solution out of the capillary back into the beaker before taking the next sample.Basic:
Allow 10.00 mL of the KMnO4- solution remaining in the buret to run into a clean 50 mL beaker. Then empty the buret of the remaining KMnO4 solution and rinse it thoroughly with distilled water. Rinse it twice, and then fill with the standard NaHSO3 solution. Titrate until the solution has turned a clear green with no trace of violet. (Just to see what happens, when the titration is finished, pour some dilute sulfuric acid into the beaker.)
Analysis:
Calculate from the volume and concentration of NaHSO3 the moles of electrons furnished by the reducing agent in each case. Then, knowing that this number of electrons is picked up by the MnO4-, calculate the moles of electrons picked up per mole of MnO4-, in each case. Assign an oxidation number to the manganese in each of the products, complete and balance each reaction according to the half-reaction method of balancing redox reactions. Explanation for Exp. 11 calculations. Calculations (.pdf)
Questions:
- Why is it that you need progressively more and more of the standard KMnO4 solution for titration as you go from acidic to neutral to basic conditions?
- Give one reason why titration of KMnO4 in acidic solution gives more precise results than does titration of KMnO4 in basic solution.
- Permangangte reacts rapidly with sulfite or bisulfite, but it does not react so fast with all reducing agents. How could a slow reaction lead to erroneous results in titrating with permanganate?
- Suppose that as you are titrating the KMnO4 in acid, some of the H2SO3 breaks down into H2O + SO2, and the SO2 leaves the solution as a gas. How will this affect your determination of the number of moles of electrons picked up per mole of MnO4-?
- There is a general rule about oxyanions that says that high oxidation states are stabilized by basic solution. Apply this rule to your observations in this experiment.
- It is found in an experiment than in an acidic solution, 20.0 mL of 0.300 M H2SO3 solution reacts completely with 10.0 mL of 0.200 M Cr2O72- solution. If the H2SO3 goes to HSO4-, what is the final oxidation state of the chromium ion?