7/28/2019 LabFePillRedox

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Redox Titration: Standardization of KMnO4 and the % Fe2+ in Iron Pills

Objectives: Standardize a stock KMnO4(aq) solution Determine the percent Fe in iron pills

A chemical analysis that is performed primarily with the aid of volumetric glassware (e.g., pipets, burets,volumetric flasks) is called volumetric analysis. For a volumetric analysis procedure, a known quantity of one

substance reacts with a to-be-determined amount of another substance. The known quantity is either aprimarystandard(a substance that has a known high degree of purity, a relatively large molar mass, is nonhygroscopic,and reacts in a predictable way) or a standard solution (a solution having a very well known concentration of asolute).

Redox titration is a process used to determine the concentration of an ion in an unknown solution by reacting itwith another ion in a solution having a known concentration. The equivalence point is reached when the totalnumber of electrons lost in the oxidation reaction is equal to the total number of electrons gained in thereduction reaction.

In this experiment, you will first standardize the approximately 0.01Mpurple-colored potassium permanganatesolution by redox titration with a primary standard, iron(II) ammonium sulfate hexahydrate. Permanganate

(MnO41-

) is a strong oxidizing agent which causes the iron to be oxidized to Fe

3+ions. The manganese is

reduced from a 7+ oxidation state in the permanganate ion to form colorless Mn2+ions. The equivalence pointis indicated at the point when all of the Fe2+ions in solution are oxidized and the colorless mixture retains apurple tint. (This color may be more orange in appearance, depending upon the concentration of the Fe2+(aq),which has a yellow-green tint. A few drops of concentrated phosphoric acid can be added to form a complexwith the iron and minimize this color.) Sulfuric acid is added to increase the concentration of hydrogen ions inthe solution.

Once you have determined the concentration of the standard solution of KMnO4(aq), you will use it to determinethe concentration of Fe2+ in iron pills.

Materials

balance50 mL buret125-mL erlenmenyer flasksmortar and pestlestock KMnO4(aq), ~0.01M

Fe(NH4)2(SO4)26H2O3MH2SO4conc H3PO4Fe pills

CAUTION: Potassium permanganate and sulfuric acid can cause chemical burns. The KMnO4 will stain skinand clothing. Avoid skin contact with these chemicals.

Standardization of KMnO4

Procedure:

1. Fill a clean, preferably glass, 50-mL buret with the ~0.0100 M KMnO4. [Standard operating procedure, orSOP, is to wash the buret with soap, rinse with tap water, rinse with distilled water, then a final rinse withtwo small portions of titrant to ensure that the titrant is not diluted with rinse water.]

2. Obtain two clean Erlenmeyer flasks; label the flasks "1" and "2". Weigh out approximately 0.5 g of iron

(II) ammonium sulfate hexahydrate, Fe(NH4)2(SO4)26H2O (abbreviated FAS) onto a piece of weighing

paper. Tare Flask 1 and tap in the FAS; record the mass to at least 0.01 g. Repeat with Flask 2. Add 10mL of distilled water, 5 mL of 3M H2SO4, and a few drops of concentrated H3PO4(aq) to each flask andswirl to dissolve the FAS.

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3. For each lab bench (2 teams): place about 50 mL of water in a beaker and add 1 drop of the permanganatesolution. This is the color standard for the reaction. When the equivalence point is reached, the colorintensity of the mixture should match this standard and remain for 5 seconds or more.

4. Record the initial volume reading in the buret to 0.05 mL. Add the permanganate to the FAS solution inFlask 1 until the equivalence point is reached. Record the final volume reading in the buret and determinethe exact volume of permanganate solution used.

5. Repeat the titration process using the FAS in Flask 2. Wash the mixtures in the flasks down the drain, then

rinse the flasks with distilled (or tap) water. From the mass of each FAS sample, calculate the moles ofFe2+ ions present. Use the balanced equation to determine the moles of KMnO4 needed to reach theequivalence point. From the volume of KMnO4 used to titrate the samples, calculate the two molarities ofyour permanganate solutions and the average molarity.

6. Add your answer to the class data on the board and wait for the teachers ok before proceeding.

Analysis of the Iron PillProcedure:

1. Grind up two iron pills in a clean, dry mortar and pestle. Weigh ~0.3 g of the powder on weighing paper;tare a clean Erlenmeyer flask, transfer the powder into the flask and mass to 0.01 g. Repeat with a secondsample. Add 25 mL of distilled water, 15 mL of 3M H2SO4, and a few drops of concentrated H3PO4(aq) toeach flask and swirl to dissolve the iron pill.

2. Record the initial volume reading in the buret to 0.05 mL. Add the permanganate to the iron pill solutionin flask 1 until the equivalence point is reached. Record the final volume reading in the buret and determinethe exact volume of permanganate solution used.

3. Repeat the titration process using the iron pill in flask 2.4. Wash the mixtures in the flasks down the drain, then rinse the flasks with distilled water. Rinse the buret

with tap water and return to the teacher.

Results and Conclusions:

1. From the volume of KMnO4 used to titrate the samples, determine the moles of KMnO4 used. Calculate themoles of Fe2+present in the pill samples based on the balanced redox equation. Determine the g Fe2+

equivalent to this number of moles, then calculate the two trials for mass % of Fe2+ in the iron pill samplesand the average % by mass composition.

2. Record your data for the experiment on the board and compare your results to the class averages. How dothe results compare with the manufacturers stated value of %Fe? Give possible reasons for anydiscrepancy greater than 5%.