Experiment #8. Redox Titration

Transcription

1 Experiment #8. Redox Titration Goal To determine the mass of iron in supplement pill using redox titration. Introduction Oxidationreduction reactions (also known as redox reactions) are reactions that usually involve transfer of electrons. To determine the number of electrons transferred, oxidation states are assigned. Oxidation states of atoms are numbers that help chemists keep track of electrons during a reaction. Each atom in an equation can be assigned an oxidation state according to certain rules. If the oxidation state of an atom increases as you go from the reactants to the products in an equation, oxidation has occurred (electrons have been lost); if the oxidation state decreases, reduction has occurred (electrons have been gained). In balancing redox reactions, the reaction is often broken down into halfreactions the reduction halfreaction and the oxidation halfreaction. For example: (1) MnO 4 + 8H + + 5e Mn H 2 O (2) Fe 2+ Fe e In the reduction halfreaction (1), manganese has undergone a decrease in oxidation state from +7 to +2. Thus each manganese atom has gained 5 electrons. In the oxidation halfreaction (2), each iron atom has undergone an increase in oxidation state from +2 to +3 that is, each iron atom has lost 1 electron. (Spectator ions have been left out in these equations.) Oxidation must occur along with reduction, and the atoms that gain or lose electrons are called the "redox pair". In a redox pair, one element will lose electrons and the other element will gain electrons. The total number of electrons lost and gained in the reaction must be equal. Therefore we must multiply the reduction half reaction or the oxidation half reaction (or both) by a small whole number to balance the reaction. In the case above, we multiply the oxidation half reaction (2) by 5. Once the above reactions are balanced, the overall redox reaction becomes: (3) MnO Fe H + Mn Fe H 2 O In Part A of the experiment today, you will find the exact concentration of a KMnO 4 solution by titrating it against a solution of iron (II) ammonium sulfate hexahydrate (also known by its former name of ferrous ammonium sulfate), Fe(NH 4 ) 2 (SO 4 ) 2 6H 2 O where the number of moles of Fe(NH 4 ) 2 (SO 4 ) 2 6H 2 O is precisely known. This process of finding the molarity of a solution is known as standardizing the solution. Using a buret, you will slowly add the KMnO 4 solution to the iron(ii) ammonium sulfate solution. Reaction (3) will occur. Permanganate ion (MnO 4 ) has a deep violet color and acts as its own indicator. As the titration proceeds, the MnO 4 (added from the buret) reacts with the Fe 2+ in the solution present in the flask, converting the MnO 4 to Mn +2. When there is no more Fe 2+ to react with, the next drop of MnO 4 that is added remains in the solution giving the solution a faint pink color. This is known as the "end point" or equivalence point of the titration and indicates the end of the redox titration. In Part B of the experiment you will use your standardized KMnO 4 solution to determine the milligrams of iron in an iron (Fe 2+ ) supplement pill. You will crush and dissolve an iron tablet, then perform a titration similar to the one in part I to determine the amount of iron in it.

2 Laboratory Activity Equipment buret mortar and pestle analytical balance buret clamp funnel ring stand 3 x 250 ml Erlenmeyer flasks Chemicals Fe(NH 4 ) 2 (SO 4 ) 2 6H 2 O KMnO 4 solution 3 M H 2 SO 4 concentrated H 3 PO 4 iron tablet deionized water Procedure A. Standardizing the KMnO 4 solution 1. Rinse a buret 3 times with deionized water; be sure to let some run through the tip. 2. Rinse the buret carefully with a small portion (2 ml or less) of KMnO 4 (aq). (Use the plastic pipettes with markings to measure 2 ml) Dispose of the used KMnO 4 (aq) down the sink. Next, fill the buret with the KMnO 4 (aq) to 0.00 ml. Be sure that the tip is filled. 3. Weigh two ~0.3 g samples of Fe(NH 4 ) 2 (SO 4 ) 2 6H 2 O on an analytical balance and record the exact masses on the data sheet. Place each sample in a 250 ml Erlenmeyer flask (use two flasks). Be sure to get the entire sample into the flask. (You can use some deionized water to rinse out the weigh dish into your flask.) Dissolve each sample in ~ 50 ml of deionized water. Using graduated pipettes (not graduated cylinders), add 3.0 ml of 3 M H 2 SO 4, and 2.0 ml concentrated H 3 PO 4 (85%) to each flask. The ions in the acids form a colorless complex with Fe +3 in solution which simplifies the detection of the endpoint. 4. Record the initial reading on the buret (read the bottom of the meniscus at eyelevel) to two decimal places. Start to add the KMnO 4 solution. At first, the violet drops of the permanganate solution will disappear quickly in the flask. When the disappearance of the violet color starts slowing down, 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 endpoint has been reached. A white piece of paper under the Erlenmeyer flask will aid in detecting color changes. Record the level of KMnO 4 solution in the buret to the nearest 0.01mL. 5. Repeat the titration in step 4 with the second sample of Fe(NH 4 ) 2 (SO 4 ) 2 6H 2 O. 6. Using the balanced equation, calculate the moles of Fe(NH 4 ) 2 (SO 4 ) 2 6H 2 O and the molarity of the KMnO 4 solution using the data from each titration. Your molarities for KMnO 4 should agree within 4%. If they don't, do a third titration. Average your molarities. B. Determining the Mass of Iron in a Pill 7. Obtain one iron supplement pill and a mortar and pestle. Grind the pill into a fine powder then transfer all of the powder to an Erlenmeyer flask. Use deionized water to rinse the contents of the mortar into your flask. Add 75 ml of deionized water, 5.0 ml of 3 M H 2 SO 4 and 3.0 ml concentrated H 3 PO 4 to the flask. Use the plastic pipettes to measure the acids. Pills contain binders and other ingredients that may not dissolve. 8. Titrate the solution containing the iron tablet with the now standardized KMnO 4 solution. The procedure is the same as in the above titrations. Again, towards the end of the titration, the KMnO 4 has to be added one drop at a time. It may take a long time to get a persistent pink color because the iron in the tablet dissolves into the solution slowly. Grinding your iron pill thoroughly in the mortar and pestle will assist the reaction to proceed as rapidly as possible. (Note: the initial color of the solution will start out as the color of the pill coating. The color will fade and will eventually become light purple at the equivalence point)

3 9. Calculate the moles of Fe 2+ present in the unknown sample and the milligrams of iron in the tablet. 10. Calculate your percent error according to the equation: % Error = mg Fe in pill mg Fe determined by titration x 100 mg Fe in pill 11. Drain any unused KMnO 4 (aq) back into the reagent bottle. Rinse your buret several times with deionized water. Be sure to rinse out the tip. Calculations A. Calculating the exact molarity of KMnO 4 Knowing the mass and molar mass of iron (II) ammonium sulfate, you can determine the moles of iron (II) ammonium sulfate. There is one mole of Fe 2+ ions per mole of iron (II) ammonium sulfate. Using the balanced equation (3), you can determine the moles of MnO 4 reacted based on the reaction stoichiometry (mole ratio) and the moles of iron used. The molarity of KMnO 4 can be determined by dividing the moles of MnO 4 by the volume of added in liters. This process of finding the exact concentration (molarity) is known as standardizing the KMnO 4 solution. B. Calculating the amount of Iron in the pill. Knowing the Average Molarity of the KMnO 4 (from part A) and the exact volume of KMnO 4 added, calculate the moles of KMnO 4 added. Using the balanced equation (3), you can determine the moles of Fe in the pill based on the reaction stoichiometry (mole ratio) and the moles of KMnO 4 added. The grams of Fe in the pill can be determined from the moles of Fe. Disposal contents of the reaction flasks sink. unused KMnO 4 in the buret empty into the KMnO 4 (aq) reagent bottle.

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5 CHM111 Lab Redox Titration Grading Rubric Name Team Name Criteria Points possible Points earned Lab Performance Printed lab handout and rubric was brought to lab 3 Safety and proper waste disposal procedures observed 2 Followed procedure correctly without depending too much on instructor or lab partner 3 Work space and glassware was cleaned up 1 Post Lab questions Data recorded clearly with proper units 2 Calculations are correct; work is shown in detail with units. 2 (Standardization) 2 (Fe pill) Question 1 (work shown clearly) 2 Question 2. 1 Question 3 2 Total 20 Subject to other additional penalties as per the instructor

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7 Redox Titration: Data Sheet Part A : Standardization of KMnO 4 solution Run 1 Run 2 Mass of Fe(NH 4 ) 2 (SO 4 ) 2 6H 2 O Name Show all calculations for one run below clearly and completely and include appropriate units: Molar Mass of Fe(NH 4 ) 2 (SO 4 ) 2 6H 2 O Moles of Fe(NH 4 ) 2 (SO 4 ) 2 6H 2 O Initial buret reading Final buret reading Volume of KMnO 4 added (in L) Moles KMnO 4 added Molarity of KMnO 4 solution Average molarity of KMnO 4 solution Part B : Calculating amount of iron in the tablet Initial buret reading Final buret reading Show all calculations below clearly and completely and include appropriate units. Volume KMnO 4 added (in L) Moles KMnO 4 added Moles Fe in pill grams Fe in pill calculated milligrams of Fe in pill (not 27mg use your data!) Report Page 1 of 2

8 Redox Titration: Post Lab Questions Name 1. Calculate your percent error for the amount of iron in the tablet. The actual amount of iron in the pill is 27 mg. 2. In Part B of the experiment, a student did not wait before recording the volume of KMnO 4 added after seeing a purple solution and proceeded to calculate the amount of iron. While calculating, the solution became colorless. How would this change the calculated amount of iron? Explain how it would differ from the true amount of iron in the pill. 3. Given the following redox equation: SnCl 2 (aq) + TlCl 3 (aq) SnCl4 (aq) + TlCl (aq) If it takes 8.52 ml of M SnCl 2 to titrate 15.0 ml of a TlCl 3 solution, what is the molar concentration of Tl 3+? Report Page 2 of 2