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A Spectrophotometric Study of the Permanganate–Oxalate Reaction An Analytical Laboratory Experiment Gene E. Kalbus* and Van T. Lieu Department of Chemistry and Biochemistry, California State University, Long Beach, CA 90840; *
[email protected] Lee H. Kalbus Department of Chemistry, California State University, San Bernardino, CA 92407
5C2O42− + 2MnO4− + 16H+ → 10CO2 + 2Mn2+ + 8H2O (1)
The rate of reaction can be further increased by the addition of Mn2+ at the beginning of the titration (8). A spectrophotometric method has been developed to study the permanganate–oxalate reaction. Part I of this experiment involves the quantitative analysis of an oxalate sample. An unknown oxalate is added to a measured excess of potassium permanganate, slowly bleaching out some of its color. The decreasing absorbance of the permanganate is followed and automatically recorded for 500 seconds. This experiment allows the student to photometrically and visually follow the course of the reaction, as well as determine the percent oxalate in an unknown sample. Part II of the experiment involves the investigation of the effect of sulfuric acid concentration and the effect of manganous ion catalyst on the rate of and outcome of the reaction. Part I. Spectrophotometric Analysis of an Unknown Oxalate
Beer’s Law Plot of Potassium Permanganate The wavelength of maximum absorbance is found to be 525 nm from the visible spectrum of a 4 × 10᎑4 M potassium permanganate solution. A linear relationship is obtained when absorbances (at 525 nm) are plotted versus molarities for a series of standard permanganate solutions (from 0.5 × 10᎑4 to 4 × 10᎑4 M). Beer’s law is therefore obeyed and from the slope
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of the plot, the molar extinction coefficient is found to be 2455 cm᎑1 M᎑1.
Spectrophotometric Observation of the Permanganate–Oxalate Reaction A measured excess of a standard solution of potassium permanganate in a 1-cm cell is treated with exactly 1.00 mL of an unknown oxalate solution (between 1.0 × 10᎑3 and 1.5 × 10᎑3 M) that is 1.5 M in sulfuric acid. The spectrophotometer is set at 525 nm and a time scan of decreasing absorbance versus time is immediately begun. From the resulting plot shown in Figure 1, the student can visually see the course of the oxidation–reduction reaction, the time required for the reaction to start, the initial rate of the reaction, the acceleration of the rate as the catalytic product, Mn2+, is produced, and the time required for the absorbance to level off, indicating the completion of the reaction. Oxalate Unknown Analysis The concentration of the unknown oxalate is calculated from the amount (mmoles) of the original potassium permanganate placed in the 1-cm cell, the amount of unreacted potassium permanganate after completion of the reaction, and
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The volumetric analysis of oxalate by titration with potassium permanganate is a classical experiment commonly performed by students in analytical laboratories (1–3) and a number of studies have been reported on its chemical kinetics and mechanisms (4–7). In a volumetric titration, potassium permanganate is often used as a titrant because of its high oxidation potential and because of its intense color that allows it to serve as its own indicator. However, because the reaction between the two components is much slower than is desirable for a volumetric titration, special techniques must be employed. The standard procedure is to add 90–95% of the required amount of potassium permanganate and to allow it to stand until the pink color of the permanganate disappears. The titration is then completed by titrating slowly at 50–60 ⬚C in sulfuric acid solution. The reaction is hastened by the elevated temperature and by the auto catalytic effect of the product, Mn2+, produced by the reaction:
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Time / s Figure 1. Time scan for 2.00 mL of 4.00 x 10᎑4 M KMnO4 in H2O treated with 1.00 mL of unknown oxalate in 1.5 M H2SO4.
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Figure 2. Effect of H2SO4 concentration on the permanganate–oxalate reaction. 2.00 mL of 4.00 x 10᎑4 M KMnO4 treated with 1.00 mL of 1.50 x 10᎑3 M oxalate in a final H2SO4 concentration of (a) 0.0 M, (b) 0.1 M, (c) 0.2 M, and (d) 0.5 M.
Figure 3. Effect of Mn2+ catalyst on the permanganate–oxalate reaction. 2.00 mL of 4.00 x 10᎑4 M potassium permanganate treated with 1.00 mL of 1.5 x 10᎑3 M oxalate in a final H2SO4 concentration of 0.5 M and a final Mn2+ concentration of (a) 0.0 M, (b) 1.0 x 10᎑5 M, and (c) 1.0 x 10᎑3 M.
the coefficients of reactants of the balanced equation (eq 1) for the permanganate–oxalate reaction.
ganous ion, as shown in Figure 3. At 0.0 M Mn2+ (Figure 3a), it takes approximately 150 seconds to complete the reaction, but at 1 × 10᎑5 M (Figure 3b) and 1 × 10᎑3 M Mn2+ (Figure 3c), the reactions are completed after 120 seconds and 50 seconds respectively.
Part II. Factors Affecting the Permanganate–Oxalate Reaction
Effect of Sulfuric Acid Concentration The effect of the sulfuric acid concentration on the oxidation–reduction reaction is studied by running time scans (at 525 nm) of 2.00 mL of 4.00 × 10᎑4 M permanganate treated with 1.00 mL of 1.50 × 10᎑3 M oxalate in different concentrations of sulfuric acid. The rate of reaction increases with sulfuric acid concentration, as shown in Figure 2. When no acid is present (Figure 2a) essentially no reaction occurs within the 500 second time period. For time scan plots of 0.1 M (Figure 2b) and 0.2 M (Figure 2c) sulfuric acid, the absorbance at first decreases as the permanganate is destroyed, but as the Mn2+ product is formed in a solution that is not sufficiently acidic, it reacts with the remaining permanganate to form a suspension of the precipitate, MnO2, which blocks the radiation path causing the absorbance to rise (2): 2MnO4− + 3Mn2+ + 2H2O → 5MnO2 + 4H+
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At higher sulfuric acid concentration such as 0.5 M (Figure 2d), the absorbance remains constant after the completion of the desired permanganate–oxalate reaction (eq 1). At 0.5 M sulfuric acid, the higher H+ concentration shifts the equilibrium of the reaction (eq 2) to the left and prevents the formation of MnO2.
Catalytic Effect of Manganous Ion The catalytic effect of manganous ion is studied by treating 2.00 mL of 4.00 × 10᎑4 M permanganate with 1.00 mL of 1.50 × 10᎑3 M sodium oxalate in a final sulfuric acid of 0.5 M with different added concentrations of manganous ion. The rate of reaction increases with concentration of the man-
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Hazards Sulfuric acid is corrosive and should be handled with care. Protective gloves and goggles must be worn when handling this acid. Potassium permanganate and sodium oxalate are harmful if swallowed and can be irritating to the eyes, respiratory tract and skin. Potassium permanganate is a strong oxidizing agent and can react violently with reducing agents. Conclusions The study of the permanganate–oxalate reaction is an analytical laboratory experiment that involves the application of basic analytical techniques and principles, stoichiometric calculations, and absorption spectrophotometry. In the experiment, not only is oxalate determined quantitatively by measuring its bleaching effect on a measured excess of standard potassium permanganate, but the catalytic effect of Mn2+ and the effect of the sulfuric acid concentration on the permanganate–oxalate reaction are examined. In addition, with the wide spread availability of versatile and inexpensive computer controlled spectrophotometers in colleges and universities, this experiment provides the students the opportunity to make use of the major functions of these modern instruments: wavelength scan, time scan, and quantitative analysis. This experiment could also be extended to the study of reaction rate and order of the permanganate–oxalate reaction (6, 7). Even though the reaction is very complex, this optional experiment could provide an opportunity for the student to use the technique to study some aspects of chemical kinetics that are typically not covered in analytical chemis-
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try. Finally, this technique of indirect spectroscopy could probably be extended to the study and analysis of other compounds that can be oxidized by permanganate. The experiment works well in the teaching laboratory and is well received by the students. After working with manually operated single-beam spectrophotometers, the students appreciate the ease, speed, and capabilities of modern spectrophotometers. The experiment can be completed in two three-hour lab periods. However, if the sections on the study of acid effects and catalytic effects of Mn2+ on the permanganate–oxalate reaction are excluded, the experiment can be modified and be completed in one three-hour lab period. W
Supplemental Material
Instructions for the students and notes for the instructor are available in this issue of JCE Online.
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Literature Cited 1. Skoog, D. A.; West, D. M.; Holler, F. J. Fundamentals of Analytical Chemistry, 7th ed.; Saunders College Publishing: Philadelphia, PA, 1996; p 369. 2. Blaedel, W. J.; Meloche, V. W. Elementary Quantitative Analysis, 2nd ed.; Harper and Row Publishers: New York, 1963; p 838. 3. Richardson, J. N.; Stauffer, M. T.; Henry, J. L. J. Chem. Educ. 2003, 80, 65. 4. Basolo, F.; Pearson, R. G. Mechanisms of Inorganic Reactions, 2nd ed.; Wiley; New York, 1967; p 178. 5. Cooke, D.O. Inorganic Reaction Mechanisms; The Chemical Society Monographs for Teachers, Number 33: Washington, DC, 1979. 6. Miles, B.; Nyarku, S. K. J. Chem. Educ. 1990, 67, 269. 7. Crouch, R. D. J. Chem. Educ. 1994, 71, 597. 8. Helsssen, J. J. Chem. Educ. 1974, 51, 386.
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