Redox Titration Percent iron (II) Essay

Redox Titration Percent iron (II) in a salt by permanganate titration Objective The purpose of this experiment is to determine the % Fe in Ferrous{Fe(II)} form in a ferrous salt by redox titration against a strong oxidant, potassium permanganate. Sodium oxalate is used as a standard to standardize the solution of permanganate. Background Reactions that involve the transfer of electrons from one substance (the reducing agent) to another substance (the oxidizing agent) can be used in quantitative volumetric analysis.

At the end-point of the titration the number of equivalents reduced must equal the number of equivalents oxidized, # equivalence in reduction = # equivalence in oxidation.

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Simply put, the number of electrons used in oxidation must equal the number of electrons used in reduction. The gram equivalent mass of a substance involved in an oxidation-reduction process is the number of grams reacting for 1 mole electron change. It is calculated from the molecular mass divided by the electron change per mole. Your oxidizing agent will be KMnO4 that in acid solution undergoes the reaction, MnO41- + 8 H+ + 5 e- !

Mn2+ + 4 H2O This reaction has a 5 e- change per mole of KMnO4 or gram equivalent mass equals onefifth the molecular mass.

Potassium permanganate has deep purple color and thus functions as it’s own indicator when being reacted with colorless substances. The MnO41- ion is violet and the Mn2+ produced is colorless. Therefore if the KMnO4 is the titrant, the end point is the first permanent pink color in the titration mixture. Note: It is important that permanganate titration be done in acidic solutions as it forms other products in neutral or basic solutions.

Solid KMnO4 cannot be obtained in a high state of purity so a standard solution cannot be made directly from the mass of KMnO4 and volume of solution. Moreover, KMnO4 tends to react with organic residues in water. You will make a solution of approximately the concentration needed and then standardize it using a second reagent of known concentration. Sodium oxalate, Na2C2O4, is a good reducing agent and can be obtained in reagent grade purity and so can be used as a primary standard. You can weigh out a specific mass of Na2C2O4, dissolve it in a specific volume of solution and calculate an exact concentration.

The reaction the oxalate undergoes is C2O42- ! 2CO2 + 2 eThis reaction has a 2-electron change per mole of Na2C2O4 or a gram equivalent mass equal to one-half the molecular mass. You will prepare a standard solution of Na2C2O4 and use it to standardize your KMnO4 solution. The titration is done in an H2SO4 solution. You will use your KMnO4 solution to analyze a solid sample containing ferrous ions Fe (II). Fe (II) is easily oxidized to ferric state Fe (III) Fe2+ ! Fe3+ + eThe reaction is run in the presence of H2SO4 for acid conditions.

Since salts that contain ferrous and ferric ions may impart color to aqueous solutions, a small amount of H3PO4 is added to complex the Fe3+ to a colorless form. This will result in a cleaner end point. It must also be noted that the amount of sulfuric acid added in the second step is less compared to the first step. Why? Calculation: Standardization of KMnO4 with Na2C2O4 Both these reagents are in solution so you can determine number of equivalents from normality and volume, (or molarity and volume).

At the end point, #eq. Red = # eq. Ox (NV) KMnO4 = (NV) Na2C2O4. Where N = normality If you are using the molarity, then you need the appropriate stoichiometry. In this case, when you balance the reaction for the redox process, the overall equation will be 2MnO41- + 16 H+ + 5 C2O42- ! 10 CO2 + 2 Mn2+ + 8 H2O That gives a molar stoichiometry of 2:5 for MnO41- to C2O42If you know the molarity of C2O42-, then the molarity of KMnO4 will be calculated as follows: M KMnO4 = 2/5 x [(M C2O42- x Vol. of C2O42- ) / Vol. f KMnO4] Your data provides all terms except the concentration of KMnO4, which you can calculate. Percent iron (II) in unknown sample To find percent by mass of iron (II) you need mass of Iron in the sample. Once again, it you are working with equivalence, then at the end point #eq Red = # eq Ox (NV) KMnO4 = (NV) Fe (II) If you are dealing with molar equivalence, the balanced redox equation is MnO41- + 8 H+ + 5 Fe2+ ! Mn2+ + 4 H2O + 5 Fe3+ Mass of Fe = 5 x [(Molarity of KMnO4 x Vol. KMnO4) x atomic mass of Fe] % Fe = (mass of Fe in the sample/sample mass) x 100%

Your sample mass is part of your data so you need to find the grams of iron from your titration data. You know the molarity and volume of the KMnO4 and you can determine the gram equivalent mass of iron and the atomic mass of iron. Thus you can calculate the grams for iron present in each of your samples and then the percent iron in the unknown. Equipment Burette, pipet, pipet filler, funnel, Erlenmeyer flasks, beaker, hot plate, balance Chemicals Sodium Oxalate solution, KMnO4 solution, sulfuric acid, phosphoric cid, distilled water, Iron (II) containing salt as unknown Safety Sulfuric acid is highly corrosive, be extremely careful, be extra careful during the titration of hot solutions Procedure Clean the buret and rinse it with distilled water. Then rinse with two 2 mL portions of your KMnO4 solution. Fill the buret with the KMnO4. Because of the dark color of the KMnO4 solution, you should read the top of the meniscus. Standardization of KMnO4 Clean and rinse with distilled water three Erlenmeyer flasks and a 10. 00 mL pipet.

Rinse the pipet with your Na2C2O4 solution then pipet a 10. 00 mL portion into each of the Erlenmeyer flasks. To each flask add about 50 mL-distilled water and about 15 mL of 3 M H2SO4 using a graduated cylinder for measurement. Warm the mixture to a temperature between 50 – 60 oC. Record the initial buret reading (should not be 0. 00). Add KMnO4 to the first Na2C2O4 heated sample with mixing until a permanent faint pink color remains in the titration mixture. Record the final buret reading. Repeat with the other two Na2C2O4 samples.

The heating is essential to speed up this reaction that is slow at room temperature. The reaction proceeds rapidly at warmer temperatures. It is also needed to assure that appropriate # of electrons are used in the redox process. Calculate the molarity of the KMnO4. Titration of Iron (II) sample Clean and rinse three 125 or 250 ml Erlenmeyer flasks. Using the container with your unknown and weighing by difference method, transfer between 0. 3 to 0. 5 g of unknown to each flask. Do these weighings extremely carefully on the analytical balance.

Dissolve each sample in about 20 ml distilled water and add about 5 mL of 3 M H2SO4 and 1 mL concentrated H3PO4. Heat the solution carefully to a temperature slightly above 60oC. Record the initial buret reading. Slowly titrate the first iron sample with the KMnO4 until a permanent faint pink remains. If you did not add H3PO4, the end point will be slightly difficult to detect. Record the buret reading and add a drop and check if the pink color deepens. Record the final buret reading. Repeat with the other two samples. Calculate the percent iron in each sample and report the average value.

Sample Data Sheet KMnO4 Standardization Trial 1 Volume of Na2C2O4 Initial buret reading Final buret reading Difference (mL of KMnO4) Moles of C2O42 Moles of KMnO4 used Molarity of KMnO4 Average Molarity Iron (II) unknown Unknown label/number _________________ Trial 1 Trial 2 Mass of the bottle with sample Mass of the bottle –(sample) Mass of the sample (g) Initial buret reading Final buret reading Difference (mL of KMnO4) Moles of KMnO4 added Moles of Iron (II) Mass of Iron in the sample % Iron in the sample Average ± Theoretical value (if available) % Error Trial 2 Trial 3

Trial 3 Questions 1. A 1. 874 g solid sample containing some Fe (II) is analyzed using potassium dichromate, K2Cr2O7, as the oxidizing agent in acid solution. The titration required 35. 40 mL of 0. 0173 M K2Cr2O7. The half reactions occurring are Fe2+ ! Fe3+ + e2+ Cr2O7 + 14H + 6e-! 2Cr3+ + 7H2O a. Write a balanced overall equation for the reaction occurring during the titration. b. Determine the moles of K2Cr2O7 used during the titration. c. Determine the percent iron in the sample. 2.

Consider the titration of oxalate solution with permanganate solution in which you determined the volume of oxalate from your pipet volume. For each of the following, state whether your volume of permanganate would be too high, too low, or unaffected and why. a. You did not rinse the water out of your pipet before filling it with oxalate. b. You did not dry the Erlenmeyer flask before adding the pipeted oxalate solution. c. You had the oxalate solution above the volume mark of the pipet.

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