7#=Y;K>VVV V W.W<fWWWW(W(XXzXxWY Y YB*YlVfYBY6 YBYBYlYBYBYBYBYBYBExperiment VI Determination of the Stoichiometry of a Redox Reaction pre-lab Assignment Reading: Before coming to your discussion section, read the following: 1. Sections 17-1 and 17-4 in Olmstead and Williams. 2. The remainder of this experiment in this manual. Pre-lab assignment to hand in: Prepare written responses to the following and submit your work to your instructor at the beginning of your discussion period: 1. Write balanced equations for the reaction of Fe2+ with MnO4- giving, respectively, Mn2+, Mn3+, MnO2, and MnO42- as the product. What stoichiometric ratio does each equation require for the reaction? 2. Calculate the weight of KMnO4 necessary to prepare 250 mL of 0.01000M standard KMnO4. 3. A sample of 0.5123 g of ferrous ammonium sulfate hexahydrate (F.W. 392.16 g/mol) was titrated with 14.92 mL of standard 0.01462M K2Cr2O7. What is the stoichiometric ratio of moles Fe2+ reacted with K2Cr2O7 at the endpoint? Introduction The purposes of this experiment are (1) to expand your experience with redox reactions, (2) to give you experience with a quantitative chemical technique, titration, and (3) to use a titration to determine the stoichiometry of the reaction of Fe2+ with potassium permanganate, KMnO4. From the stoichiometry of the reaction you can determine the identities of the reaction products. Titrations: A titration is a process in which a solution of known concentration is mixed with a solution of unknown concentration and a specic chemical reaction between the two reactants is carried just to completion. If the balanced equation for the reaction is known, then the concentration of the unknown reagent can be determined with excellent accuracy. The point at which the titration reaction is complete, with no excess of either reactant, is called the equivalence point. Effective titrations require accurate methods for measuring solution volumes, usually with a buret, and also the "end point" of the reaction. The "end point" of the reaction is an observable change in a measurable quantity of the reaction solution, often a color change. The "end point" (observable change) should be very close to the equivalence point at which an exactly stoichiometric amount of the known reactant has been added. In the following procedure we will use the deep purple color of the KMnO4 to advantage. When the permanganate ion reacts with Fe2+ the products are colorless. Thus, as the KMnO4 solution is added to carry out the reaction, the deep purple color is dissipated. However, as soon as the equivalence point is reached, excess purple MnOA(-,4) ion accumulates in the reaction solution and can be visually detected. The color of MnOA(-,4) ion is so intense that it can be seen at very low (~2x10-6M) concentration and our "end point" (appearance of purple color) is very close to the true equivalence point of the reaction. Redox Reactions: Potassium permanganate, KMnO4, is commonly used as an oxidizing agent for oxidation-reduction titrations. Its oxidizing properties also make KMnO4 useful in medicine as an antiseptic and an antidote for many poisons. In redox titrations, the very strong oxidation potential of KMnO4 makes it useful for the determination of a wide variety of inorganic and organic species including Sn, Fe, V, Mo, W, U, Ti, HNO2, and oxalic acid. The product of the reduction of KMnO4 depends on the reaction conditions. In our reaction, the product formed from Fe2+(aq) is Fe3+(aq) and possible products from the MnOA(-,4) ion are Mn2+(aq), Mn3+(aq), MnO2(s) or MnOA(2-,4)(aq). The reaction product and the number of electrons gained by KMnO4 must be known before using the reagent in analytical determinations. Thus you will use your titration in a slightly different manner than that described above. You will have two solutions of known concentration and will start with an unknown reaction equation. As part of the pre-lab assignment, you will write balanced equations for all of the likely reactions. Then hopefully you will be able to t your results to only one of the equations. The approach used in this experiment can be illustrated with a parallel example using perchlorate ion, ClO4-, as the oxidizing agent. In the titration of Fe2+ with ClO4-, the two possible chlorine- containing products are Cl2 and Cl-. The possible half reactions are: Fe2+ F( , )> Fe3+ + 1e- oxidation half-reaction (1) and ClO4- F( , )> 1/2 Cl2 (g) reduction half-reaction (2) or ClO4- F( , )> Cl- reduction half-reaction (3) and the balanced equations for the two possible reactions are (verify the balancing of these reactions as practice): 7 Fe2+ + ClO4- + 8 H+ F( , )> 7 Fe3+ + 1/2 Cl2 (g) + 4 H20 (4) or 8 Fe2+ + ClO4- + 8 H+ F( , )> 8 Fe3+ + Cl- + 4 H20 (5) The stoichiometry of the reaction in terms of moles Fe2+ to moles ClO4- can, therefore, be used to determine the products of the reaction: 7:1 for Cl2 as the product or 8:1 for Cl- as the product. To determine the stoichiometry, a titration was carried out. A carefully weighed sample of 0.3532 g of ferrous sulfate FeSO4.7H20 (F.W. 278.03 g mol-1) was titrated with a 0.01062M solution of KClO4. The endpoint was found at 14.99 mL of added KClO4. What is the stoichiometry of the reaction? ______________________________________________________________________________ Solution (keeping one extra significant figure and rounding at the end) The number of moles of Fe2+ is 0.3532 g of FeSO4.7H20 ( F(1 mole,278.03g)) = 1.2704 x 10-3 mol Fe2+ The number of moles of ClO4- added is 14.99 mL of KClO4 (F(1L,1000mL)) (F(0.01062mol,L)) = 1.5919 x 10-4 mol ClO4- The ratio of Fe2+ to ClO4- is F(moles Fe2+,moles ClO4-) = F(1.2704 x 10-3 mol,1.5919 x 10-4 mol) = 7.980 The stoichiometry is, therefore, 8:1 and reaction (5) is the proper reaction. ______________________________________________________________________________ The stoichiometry of 8:1 then shows the product of the reaction to be Cl- and for every mole of ClO4-, 8 electrons are transferred. Quantitative Laboratory Work: Now that you have a signicant amount of experience in the chemistry lab, you are ready to do careful quantitative determinations. To encourage you, this experiment will be graded primarily on the quality (precision and accuracy) of the results. The procedure is designed to approximate "real world" analytical work. You will prepare all of your own reagent solutions, carry out the analysis as many times as you feel is necessary and report your best estimate of the "correct" result. You should take advantage of the information in Experiment V-Appendices 1 & 2 concerning errors and discarding obviously erroneous data. Procedure Determination of the Stoichiometry and Reaction Products of the Titration of Fe2+ with MnO4- Ferrous ammonium sulfate hexahydate, Fe(NH4)2(SO4)2.6H2O will be used as the source of Fe2+, and potassium permanganate, KMnO4, will be the source of MnO4- ion. Outline: A. Prepare a 0.01M solution of KMnO4. B. Weigh 3 samples of ferrous ammonium sulfate hexahydrate. C. Titrate each sample to a faint pink endpoint. D. Determine the molar ratio of Fe2+ to MnO4- reacted at the endpoint and average the trials. Note: The weighings in parts A & B may be done the week before the actual experiment begins. However, do not dissolve any of the solid sample until you are ready to do the titrations. The solutions are stable for only about one day. If you wish to do additional trials in a subsequent week, you will need to prepare a new KMnO4 solution. A. Preparation of standard 0.01M KMnO4: Prepare an approximately 0.01M KMnO4 solution by weighing, to 4 significant figures, enough solid KMnO4 salt to prepare 250 mL of solution. Weigh into a clean, dry weighing bottle from a weighing bottle filled with KMnO4. Quantitatively transfer the KMnO4 into a 250-mL volumetric flask using a long stemmed funnel, washing the weighing bottle several times with deionized water from a wash bottle. Add approximately 100 mL of water to the ask and swirl the mixture to dissolve the solid. Dilute to the mark, adding the last few mL with a Pasteur pipet to avoid going over the mark. Shake well, while inverting the volumetric flask at least 12 times. B. Weigh three samples of ferrous ammonium sulfate hexahydrate: Weigh, to 4 significant figures, three 0.5 g samples of Fe(NH4)2(SO4)2.6H2O into clearly marked 200-mL Erlenmeyer flasks. Remember to keep track of the weight in each flask, since they will all be different. C. Titrate: Be certain that your buret is clean (no droplets adhering to the inside upon draining) and the stopcock is tightened. Rinse the buret with about 3 mL of your standard KMnO4 solution. Fill the buret using a funnel. Clear all bubbles from the tip. Remove any drops from the tip. Take an initial reading, before each titration. Because of the dark color, the bottom of the meniscus is not visible, and the top of the meniscus will be read at the beginning and end of the titration. Read the volume to two significant figures past the decimal point (for example, 1.26 mL). Add approximately 25 mL of 0.5 M sulfuric acid, H2SO4, and 25 mL of 2.0 M phosphoric acid, H3PO4, to each sample just before titrating. Place the tip of the buret well inside the neck of the Erlenmeyer flask. Titrate each sample with KMnO4 until the appearance of a very faint pink color that persists for 30 seconds. As you approach the endpoint, the pink color will begin to persist. At this point, add titrant slowly, transferring partial drops from the buret tip by touching the tip to the side of the flask. Rinse the walls of the Erlenmeyer flask after each addition with a stream of water from your wash bottle. Repeat the titration on successive samples until the precision, measured as the standard deviation from the mean of the stoichiometric ratio of Fe2+ to MnO4-, is less than 1% (i.e. less than 10 parts per thousand). D. Calculation: Calculate the concentration of your KMnO4 solution from the mass of KMnO4 you weighed out. Use the example as a guide to calculate, for each sample, the stoichiometric ratio of Fe2+ to MnO4-. Your results should have 4 significant figures. Average the trials and calculate the standard deviation from the mean (see Appendix 1). If any trial looks to be in error, use the Q test (see Appendix 2) to reject any discordant data. Report 1. Write balanced equations for the reaction of Fe2+ with MnO4- giving, respectively, Mn2+, Mn3+, MnO2, and MnO42- as the product. You did these as part of the prelab assignment, but be sure that copies of all the balanced equations are included with your report. 2. Give an example calculation of the molar ratio of Fe2+: MnO4- for one of your samples. 3. Tabulate the results for each sample that you ran (see the form below). Indicate which, if any, are not included in the mean and why (e.g. spill, Q-test). Give the average and standard deviation of the remainder. SampleFe(NH4)2(SO4)2.6H2O weightKMnO4 weight molarity of KMnO4mL KMnO4 addedmoles Fe2+ moles MnO4-1234 Average ratio Standard deviation ________________ 4. Round the average to the nearest whole number and select the proper reaction. Report the products of the reaction. 5. Answer the following question: Hydrated salts, like ferrous ammonium sulfate hexahydrate, are difficult to obtain with precisely the correct amount of hydrated water. If the salt were slightly dehydrated, that is, less than 6 moles of H2O per formula weight, what error in the final stoichiometric ratio would there be (higher, lower or no change)? Explain your answer. Extra Credit If you have time you might want to use your standardized KMnO4 solution to titrate an unknown sample. Obtain a commercial iron supplement tablet from your instructor. Grind the tablet in a mortar and pestle. Treating your sample as you did before, titrate with standard KMnO4 solution. Calculate the weight of Fe in the tablet. You will need the stoichiometric ratio determined in this experiment to calculate the number of moles of Fe2+ and hence the weight of Fe in the sample. Experiment V-Appendix 1 Random Errors and Standard Deviation If a measurement is made or an experiment repeated a large number of times, a range of values will be obtained that is due to the random error inherent in any measurement. Of the random errors, small errors are more probable than large errors and negative deviations are as likely as positive ones. The mean value of a set of measurements is the most probable value. The spread or dispersion of the results is expressed by the standard deviation, s, as follows: s = R(F(( xi - XTO(x))2,n - 1)) where xi is an experimental data value, XTO(x) is the mean and n is the number of replications. This formula actually gives only an estimate of the standard deviation unless the number of measurements is very large (in principle, an infinite number). We must recognize that when we repeat a measurement only two or three times, we are not obtaining a very large sample of measurements and the confidence that we can place in the mean value of a small number of measurements is correspondingly reduced. The following example will illustrate the procedure for calculating the standard deviation. A student determined % Cl in three different samples of unknown. The results were: 45.32%, 45.35% and 45.28%; n=3; XTO(x) = 45.31. Calculated value Deviation Square of deviation (xi) (xi XTO(x)) (xi XTO(x))2 XTO( 45.32 0.00 0.0000 ) 45.35 +0.03 0.0009 45.30 -0.02 0.0004 0.0013 s = R(F(0.0013,2)) = 0.0255 The best value for these calculations is written as 45.32 0.03% Experiment V-Appendix 2 Rejection of Discordant Data When one in a series of 3 to 10 measurements appears to deviate from the mean by more than seems reasonable, calculate the quantity Q, which is the quotient of the difference between the value under suspicion and the value in best agreement with it divided by the range of all of the values in the set. Compare the value of Q with the critical value, Qc, in Table I corresponding to the number of observations in the series. If Q > Qc, the suspect measurement should be rejected. For example, in an analysis of a compound for carbon, three trials were carried out. The results of the trials yielded three % carbon values: 36.26, 36.37 and 36.40. The last two values are nearly equivalent, but the first varies considerably from them. There was no obvious experimental error which indicated that any of the values was less likely than the others to be correct. Including the first value changes the average for the analysis significantly. Should the value 36.26 be rejected when compared to 36.37 and 36.40? To decide if the "outlying" value should be rejected, we determine its Q value and compare it to the Qc value in Table I for three determinations. Q = F(36.37 - 36.26,36.40 - 36.26) = 0.79 for N = 3, Qc = 0.94 so that Q < Qc and the datum cannot be rejected. Table I N Qc (90% confidence) 2 -- 3 0.94 4 0.76 5 0.64 6 0.56  Developed by T.W. Shattuck, Chemistry Department , Colby College, 1989. CH 141 Lab: Expt. VI -- -- CH 141 Lab: Expt. 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