Redox Lab By Maya Parks Partner: Allison Schaffer 5/21/15 Abstract: This lab was performed to determine the molarity of a solution of iron (II) sulfate by titrating it with a.0200m solution of potassium permanganate. We were able to determine the molarity to be 0.229M, although in reality, it was 0.200M. This was perhaps because we over titrated the solution slightly, causing the number to be a little too high. Also, we had water in our graduated cylinder when measuring it out, which could cause it to be lower, and we should have cleaned our equipment. Our lab had on average 15% error due to these mistakes. However, this lab was successful as it demonstrated the uses of redox reactions, and how to apply the concepts we learned to a real world example of a titration. Purpose: We will be observing a redox reaction and using the molarity of one solution in order to determine the molarity of the other, in effect, performing a titration. We will apply concepts of redox in this experiment as well as from previous units, such as those in the stoichiometry unit. We will determine the half reactions, and apply a fundamental understanding of redox reactions, in an effort to find the molarity of the solution of iron (II) sulfate. Materials: 0.020M KMnO 4 15.0mL 3.0M H 2 SO 4 30.0mL unknown molarity FeSO 4 Burette Burette clamp Ring Stand Erlenmeyer Flask Distilled Water Beaker
Graduated Cylinder Goggles Procedure Condition the buret with about 5 ml of the 0.020 M KMnO 4 solution and discard rinse down the drain. Fill the buret with potassium permanganate. Put 50 ml of distilled water into a clean erlenmeyer flask and add one drop of potassium permanganate. This solution will be used as a color standard. Measure 10.0 ml of the iron (II) sulfate solution into a clean erlenmeyer flask Add about 5 ml of 3.0M sulfuric acid to the flask and immediately begin titrating with the potassium permanganate solution. When the endpoint is reached read the buret and record the final readings. Rinse the flask with distilled water. Repeat the entire procedure (steps 6-8) two more times. Rinse the solutions down the drain with plenty of water. Data: Trial 1 2 3 FeSO 4 in ml 9.38 9.85 9.76 Initial volume 0.00 22.00 0.00 Final volume 21.75 44.50 22.00 Calculations: 1. Calculate the volume of KMnO 4 used for each trial a. Subtract initial from final volume V 2 = 22.50 ml V 3 = 22.00 ml 2. Calculate the average volume of KMnO 4 used
a. Add volumes from 1. And divide by 3 3. Using the average volume you calculated in step 2, calculate the number of moles of MnO 4 - used a. Convert ml into L and then multiply by the molarity to get moles 4. Write the half reactions for the REDOX reaction in this lab a. Mn +7 + 5e - Mn +2 Fe +2 Fe +3 e - 5. Write the overall balanced net ionic equation. a. 8H + MnO 4 - + 5Fe +2 Mn +2 + 5Fe +3 + 4H 2 O 6. Using the molar ratios of the reactants, calculate the number of moles of Fe 2+ used. a. Take moles of MnO 4 - and multiply by the molar ratio 7. Calculate the Molarity of the FeSO 4 Solution.
Error Analysis: In our experiment, we over titrated the solution just a bit too much. This caused us to have a greater volume of permanganate, which caused us to calculate a greater number of moles of FeSO 4 and the molarity was greater than it should have been. Additionally, there was often water still left in the graduated cylinder when we would measure out 10mL of iron (II) sulfate, which meant there was more water and a smaller number of moles of FeSO 4 in the flask when titrating so we would have used a smaller number of moles to titrate and when we calculated it, we would have ended up with a lesser concentration. It would be very helpful to go slower when titrating the solution, so as to eliminate the error of over titrating. Also, it would be more efficient to dry the equipment, or perhaps condition it before measuring out the solution, like we condition the burette. Conclusions: 1. During the REDOX reaction the MnO 4 - ion is changed to Mn 2+. How many electrons are involved? Is it an oxidation or a reduction? Explain your answers? 5 electrons are gained by the Mn ion, showing that this is a reduction. The acronym OIL RIG shows that reduction is the gaining of electrons as the positive charge drops. 2. What is the purpose of the solution prepared in step 4 of the procedure? The solution provides a comparison for when the solution is fully titrated. 3. Why is sulfuric acid added to the titration flask in step 6 of the procedure? The sulfuric acid provided an acidic environment for the reaction allowing there to be loose H + ions and water molecules that participate in the reaction. 4. If you titrated a second unknown solution of iron (II) sulfate, and less potassium permanganate was required to reach the endpoint, what could you conclude about the concentration of the second solution compared to the first one. Explain your answer. The concentration of the second unknown solution is less than the first solution. 5. If you did not condition your buret after rinsing with distilled water, how could this affect your calculated final concentration? Explain your answer.
The concentration of the solution in the burette would be less as part of it would be water. A greater volume would therefore be required to titrate the solution, and the calculated concentration of the unknown solution would be greater than in reality 6. A 0.345 gram sample of anhydrous BeC 2 O 4, which contains an inert impurity, was dissolved in sufficient acidic water to produce 100.mL of solution. A 20 ml portion of the solution was titrated with KMnO 4 (aq). The unbalanced equation for the reaction that occurred is as follows: MnO 4 - (aq) + C 2 O 4 2- (aq) Mn 2+ (aq) + CO 2 (g) The volume of 0.0150M KMnO4(aq) required to reach the equivalence point was 17.80mL. a) Write the half reactions for this reaction Mn +7 + 5e - Mn +2 C 2 +3 2C +4 2e - Net Ionic: 2Mn +7 +5C 2 +3 2Mn +2 + 10C +4 b) Identify the reducing agent in the titration reaction C 2 O 4 2- c) For the titration at the equivalence point, calculate the number of moles of each of the following that reacted i. MnO 4 - (aq)= ii. C 2 O 4 2- (aq)= d) Calculate the total number of moles of C 2 O 4 2- (aq) that were present in the 100. ml of prepared solution. e) Calculate the mass percent of BeC 2 O 4 (s) in the 0.345 gram sample.