CHEMISTRY 221 LABORATORY—QUANTITATIVE ANALYSIS
EXPERIMENT: PERMANGANIMETRIC DETERMINATION OF IRON IN IRON OXIDE
Adapted from J. Chem. Ed., 1992, 69, 935.
Potassium permanganate has been widely used as an oxidizing agent for over 100 years. It is a reagent that is readily available, inexpensive, and requires no indicator unless very dilute solutions are used. One drop of 0.1 N permanganate imparts a perceptible pink color to the volume of solution normally used in a titration. This color is used to indicate excess of the reagent. Permanganate undergoes a variety of chemical reactions, since manganese can exist in oxidation states of +2, +3, +4, +6, and +7.
The most common reaction encountered
in the introductory laboratory is the one that takes place in very acidic
solutions, 0.1 M or greater:
MnO4- + 8 H+ + 5 e- Mn+2 + 4 H2O Eº = +1.51 V
Special precautions must be taken in the preparation of permanganate solutions. Manganese dioxide catalyzes the decomposition of permanganate solutions. Traces of MnO2 initially present in the permanganate, or formed by the reaction of permanganate with traces of reducing agents in the water, leads to decomposition. Directions usually call for dissolving the crystals, heating to destroy reducible substances, and filtering through asbestos or sintered glass (non-reducing filters) to remove MnO2. The solution is then standardized, and if kept in the dark and not acidified, its concentration will not change appreciably over a period of several months.
of Permanganate Solutions: Sodium Oxalate
Sodium oxalate, Na2C2O4, is a good primary standard for permanganate in acid solution. It can be obtained in a high degree of purity, is stable on drying, and is non-hygroscopic. Its reaction with permanganate is somewhat complex, and even though many investigations have been made, the exact mechanism is not clear. The reaction is slow at room temperature, and hence the solution is normally heated to about 60 ºC. Even at an elevated temperature the reaction starts slowly, but the rate increases as manganese(II) ion is formed. Manganese(II) acts as a catalyst, and the reaction is termed autocatalytic, since the catalyst is produced in the reaction itself. The equation for the reaction between oxalate and permanganate is
5 C2O4-2 + 2 MnO4- + 16 H+ 2 Mn2+ + 10 CO2 + 8 H2O
For a number of years analysts employed the procedure recommended by McBride, which called for the entire titration to be carried out slowly at elevated temperature with vigorous stirring. Later, Fowler and Bright thoroughly investigated the reaction and recommended that almost all of the permanganate be added rapidly to the acidified solution at room temperature. After the reaction is complete, the solution is heated to 60 ºC and the titration completed at this temperature. This procedure eliminates any error caused by the formation of hydrogen peroxide.
Analytical Procedure: Iron in
The determination of iron in iron ores is one of the most important applications of permanganate titrations. Before titration with permanganate any iron(III) must be reduced to iron(II). This reduction can be done with the Jones reductor (a column consisting off Zn granules amalgamated with Hg or with tin(II) chloride. The Jones reductor is preferred if the acid present is sulfuric, since no chloride ion is introduced. If Sn(II) is used, a slight excess of Sn(II) chloride is added to ensure completeness of reduction. This excess must be removed or it will react with the permanganate upon titration. For this purpose, the solution is cooled, and mercury(II) chloride is added rapidly to oxidize excess tin(II) ion. Iron(II) is not oxidized by the mercury(II) chloride. The precipitate of mercury(I) chloride, if small, does not interfere in the subsequent titration. However, if too large an excess of tin(II) chloride is added, mercury(I) chloride may be further reduced to free mercury:
Mercury, which is produced in a finely divided state under these conditions, causes the precipitate to appear gray to black. If the precipitate is dark, the sample should be discarded, since mercury, in the finely divided state, will be oxidized during the titration. Tin(II) chloride is normally used for reduction of iron in samples which have been dissolved in hydrochloric acid.
A method in which zinc is used as the reductant in situ rather than as a Jones reductor, has several advantages over conventional determination of iron. It does not require the use of mercury(II) compounds, which are environmentally problematic and it avoids the preparation and upkeep of Jones reductors. A desirable feature of this procedure is the residues from the titration, Mn+2, Fe+3, Zn+2, Na+, Cl-, PO43-, and SO42- are relatively innocuous once the excess acid is neutralized. A minor problem with the reduction of iron by zinc is the large excess of HCl required to dissolve the ore sample. To minimize the amount of zinc consumed by the HCl, some of the excess HCl is removed by boiling briefly after dissolution of the ore is complete.
A solution of manganese(II) sulfate, sulfuric acid, and phosphoric acid, called "preventive," or Zimmermann-Reinhardt solution, can be added to the hydrochloric acid solution of iron before titration with permanganate. Phosphoric acid complexes with Fe(III) to lower the concentration of free Fe3+ helping to drive the titration reaction to completion. Also, the iron(III)-phosphate complex is colorless which helps in reducing the yellow color that Fe3+ exhibits in chloride solutions.
The reaction for the titration of Fe+2 by MnO4-
MnO4- + 8 H+ + 5 Fe2+ Mn+2 + 5 Fe3+ + 4 H2O
|zinc metal foil||3 M H2SO4|
|conc. HCl||conc. H2SO4|
|potassium permanganate||sodium oxalate|
Preparation of 0.02 M ( 0.1 N ) Potassium Permanganate
Use the top loading balance to weigh approximately 3.2 g of reagent grade KMnO4 into a clean 1.0 L amber glass bottle. Add enough distilled water to fill the bottle to approximately 1/3 full. Place on a magnetic stirrer and stir for at least 10 minutes. Finish filling the bottle with distilled water. Continue stirring until you are ready to titrate. Any inhomogeneity of the titrant must be avoided.
Standardization of Potassium Permanganate - Fowler-Bright Method:
Weigh accurately three sample of 0.25-0.30 g each of dried sodium oxalate into clean 500-mL Erlenmeyer flasks. Dissolve in ~50 mL of distilled water.
Add an additional 200 mL of water followed by 30 mL of 3 M H2SO4. Insure that the solid is completely dissolved.
Steadily add permanganate titrant directly to the oxalate solution (not down the walls of the flask) and stir slowly.
Add sufficient permanganate (35 to 40 mL) to come within a few milliliters of the equivalence point, add at a rate of 25 to 35 mL/min.
Let the solution stand until the pink color disappears. (This may take 30 to 45 seconds since the reaction is not instantaneous. If the color does not disappear after about 5 minutes, indicating excess permanganate , discard the solution and add less permanganate to the next sample.)
Heat the solution to 55 to 60 °C.
Complete the titration at this temperature, adding permanganate slowly until one drop imparts to the solution a faint pink color that persists for at least 30 seconds.
Add the last milliliter slowly, allowing each drop to be decolorized before adding another. The solution should be as warm as 55 °C at the end of the titration.
Prepare a blank by adding 250 mL of water and 50 mL of 3 M H2SO4 and heating to 55 to 60 °C.
Add permanganate solution dropwise until the color matches that of the titrated solution.
Subtract this volume (usually about 0.03 to 0.05 mL) from the volume used in the titration.
After the titration of three samples, calculate the molarity obtained in each titration and average the results. With care, the average deviation should be as small as ±0.2%.
Calculate the concentration of the KMnO4 solution at this point:
Iron Ore Analysis:
In a hood, add 15 mL of concentrated HCl to accurately weighed samples of unknown, approximately 0.4 g in 500-mL Erlenmeyer flasks. (Avoid spillage of the unknown as it is very hard to clean up.)
Cover with a ribbed watch glass (or an ordinary watch glass with glass hooks). Heat nearly to boiling in the hood (a large hotplate is convenient), swirling occasionally until the ore is dissolved, about 5-10 min. Some white residue of silica may remain and can be ignored. Add additional HCl only if evaporation reduces the volume significantly; that is, iron oxide is reprecipitated.
Following dissolution, boil the solution gently for a few minutes to expel as much HCl(g) as possible. (The HCl concentration cannot drop much below 6 M, but that does reduce somewhat the amount of zinc required for reduction.)
Cool briefly, and add about 1 g of zinc metal in 3 portions, replacing the watch glass immediately. (Zinc ribbon reacts more readily and is easier to dispense than granular zinc, 1 g being about 3-4 cm of 0.0005-in. ribbon. If using granular zinc, 20 mesh appears to react well. Do not add all of the zinc at once since this may reduce the Fe3+ to metallic iron.)
Heat the flask as necessary while swirling it gently, to maintain a vigorous but not violent reaction, until the zinc is consumed. A few flecks of non-reactive black residue may be ignored.
Add a second 1 g of zinc in 3 portions and heat the solution as necessary.
Repeat the addition zinc in smaller portions until reduction is complete. Usually this means that the solution is clear with no trace of yellow. This reaction is, in essence, a titration of the Fe3+ with Zn0. Adding too much extra zinc metal will greatly increase the time required to dissolve all of the excess zinc. (In concentrated Fe2+ solutions, a very light green hue of iron(II) may be apparent in a solution that is completely reduced; a solution that appears green with a trace of yellow will require a small additional portion of zinc. In addition, a solution that is too low in volume may appear to show a yellow-green color — add some distilled water and see if the color of the solution turns clear, a sign that all of the iron(III) has been reduced.)
When the solution appears to be completely reduced, add another 0.1 g of zinc and heat nearly to boiling for about 5 min. (Approximately 3 g of zinc will suffice for reduction of 0.4 g samples dissolved as described above. The final portion of zinc should not be added until one is ready to complete that particular titration, since iron(II) undergoes air oxidation.)
If any zinc remains undissolved, add 5 mL of concentrated H2SO4 in 20 mL of water, and swirl the flask until the remaining zinc is dissolved. Very fine bubbles of H2 will emanate from any undissolved zinc as opposed to the larger bubbles of a boiling solution. Set your reaction mixture off the hotplate and allow the boiling to stop to check for any "fizzing" unreacted zinc.
Add 15 mL of Zimmermann-Reinhardt reagent (79 g MnSO4, 100 mL of concentrated H2SO4, and 200 mL of concentrated H3PO4, diluted to 1 L) and dilute the solution to about 150 mL with water. (If no water has been added to the solution, reduction is rapid, taking a total of 10-15 min, and the higher concentration of iron makes it easier to see the color change that indicates complete reduction. A solution which has been diluted with water may take considerably longer to reduce. The greenish tint of the solution, which intensifies as the titration progresses, makes endpoint detection a bit troublesome but not particularly difficult. Dilution with water minimizes this side reaction and makes the endpoint much easier to see. The permanganate endpoint thus appears as a change from pale green to tan to a persistence of the pink color. An absolutely white background for the flask is of considerable help in pinpointing the endpoint, and an initial "trial" is invaluable for most students.)
Titrate immediately with standard 0.02 M KMnO4.
After completing the titration, carefully neutralize the contents of the flask (using 50% NaOH) before discarding into waste containers or as directed by the instructor.
Calculate the percentage of iron in your unknown at this point:
Report the result in terms of the percentage of iron in your unknown. The range of unknown values should be 20% to 70%.
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