Exp 3 Oxidation-Reduction Determination of Iron by Titration with Potassium Permanganate-1 PDF

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Oxidation-Reduction lab report...

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CHEM-102

Experiment #3

Oxidation-Reduction: Determination of Iron by Titration with Potassium Permanganate Pre-Lab Activities: 

After reading the lab, complete items a, b, c, and d (title, purpose, chemicals and equipment, and summary of procedure) as described on page 10 of Exp. 1 on an 8 1/2 x 11 sheet of paper.



Answer the following questions on 8 1/2 x 11 sheet of paper or in your laboratory notebook if one is required by your instructor:

1. Write the balanced net ionic equation for the reaction between MnO4- ion and Fe2+ ion in acid solution. 2. How many moles of Fe2+ ion can be oxidized by 1.2 x 10 -2 moles MnO4- ion in the reaction in Question 1? 3. A solid sample containing some Fe2+ ion weighs 1.923 g. It requires 36.44 mL 0.0244 M KMnO4 to titrate the Fe2+ in the dissolved sample to a pink end point. a. How many moles MnO4- ion are required? b. How many moles Fe2+ are there in the sample? c. How many grams of iron are there in the sample? d. What is the percentage of Fe in the sample? Lab Activities: 

Go over the sample calculations and prelab questions with your lab instructor



Complete lab and fill in data sheet for the Standardization of Potassium Permanganate. The next data sheet, Percent Iron in Unknown Sample will be completed the following lab period.

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Objective:



To illustrate the concept of oxidation-reduction reactions and to use this reaction to determine the amount of iron in an unknown sample.

Introduction: Oxidation is defined in the broadest sense as a reaction in which the oxidation state of an atom or ion becomes more positive, and reduction as a reaction in which the oxidation state of an atom or ion becomes more negative. Thus, these two processes must always occur simultaneously and to the same extent in a reaction. In the reaction of sodium with chlorine to form sodium chloride, the oxidation state of sodium changes from 0 to +1, while that of chlorine changes from 0 to -1. Sodium undergoes oxidation and chlorine undergoes reduction. 2 Nao + Cl2o  2 Na+1 + 2 Cl-1 Chemical reactions involving oxidation and reduction, or redox reactions as they are called collectively, have many analytical applications. Redox titrations are more widely used than all other titration methods combined. Although many substances undergo redox reactions, most important analytical applications are accomplished with a limited number of reagents. Some of the more important reagents are potassium permanganate, potassium dichromate, potassium iodate, and the cerium (IV) ion. The endpoint for a redox titration may be detected by several methods. The color of the reagent itself may be intense enough to serve as an indicator. For example, the purple color of the potassium permanganate can be used in a number of reactions. Some methods employ highly colored organic molecules that change colors during the course of the reaction. Other methods utilize a measurement of the electrical potential of the solution to detect the endpoint. Potassium permanganate (KMnO4) has been selected for use in this experiment, the titration of iron in an unknown containing iron(II) ammonium sulfate, Fe(NH 4)2(SO4)2 • 6H2O. In acid solution the MnO4- (Mn is in oxidation state of +7) undergoes reduction to Mn2+ (oxidation state of +2) according to the equation +7

+2

8 H+ + MnO4- + 5 e-  Mn2+ + 4 H2O The Mn must gain five electrons per atom. The iron undergoes oxidation from Fe2+ to Fe3+, going from an oxidation state of +2 to +3. Each iron atom loses one electron. Fe2+  Fe3+ + e-

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Obviously, it would take five Fe atoms for each Mn atom in order for the number of electrons lost to be equal to the number of electrons gained. The balanced equation for the reaction is: 8 H+ + MnO4- + 5 Fe2+  Mn2+ + 4 H2O + 5 Fe3+ The MnO4- ion is purple (pink in dilute solution) and the Mn 2+ ion is nearly colorless. The appearance of the first permanent pink color due to a slight excess of MnO4- ion signals the end of the titration. The presence of phosphoric acid sharpens the endpoint. Procedure for First Lab Period (Standardization of KMnO4): WEAR YOUR SAFETY GLASSES WHILE PERFORMING THIS EXPERIMENT Before beginning the titration to find the percent of iron in ferrous ammonium sulfate (second lab period), standardization must be performed on the KMnO4 solution. Standardization will determine the exact molarity of the solution for later use in the titration of the unknown iron sample. The potassium permanganate solution can be standardized by performing the following procedure: 1. Clean three 125 ml or 250 ml Erlenmeyer flasks. 2. Clean a 25 ml burette and rinse it with deionized water. Fill the burette with two or three milliliters of KMnO4 solution and swirl it throughout the entire burette. Pour the solution out of the burette into the inorganic liquid waste container and repeat a second time. 3. Fill the burette with the KMnO4 solution until the meniscus is at the zero mark on the burette. 4. Weigh three portions of approximately 0.5 g each (record the mass to the 4th decimal!) from the same sample with known iron percent (make sure to record the percentage of iron in your lab notebook/data sheet) and put each portion into a clean 125 ml or 250 ml Erlenmeyer flask. 5. Add 25 ml of 1 M sulfuric acid (dilute H2SO4) to each Erlenmeyer flask and swirl until crystals completely dissolve. Add 2 mL of 85% phosphoric acid and swirl. 6. Immediately begin the titration with KMnO4. 7. The endpoint is reached when a faint pink color persists for at least 30 seconds. Record the amount of milliliters of KMnO4 solution required to reach the endpoint. 8. Repeat steps 6-8 for the second and third samples. Using the known percent of iron, find the molarity of KMnO4.

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Procedure for Second Lab Period (Titration of Unknown Percent Iron Sample): 1. Obtain an unknown iron sample from your instructor. Accurately weigh out approximately 0.5 gram of your unknown sample (record the mass to the 4th decimal!). Carefully transfer the measured sample to a clean, numbered 250 mL flask. Repeat this procedure for two additional portions of the same unknown sample and record their masses to the 4th decimal. 2. Clean a 25-mL burette thoroughly and place it in a burette clamp fitted to a ring stand. Fill the burette with the KMnO4 solution, making sure that the tip of the burette is free of air bubbles. Drain the KMnO4 solution to just below zero and record the initial reading in the space provided for sample #1. Record the molarity as determined in the standardization procedure from the previous lab session. 3. To unknown iron sample #1 add 25 mL of 1 M H2SO4 and dissolve the sample completely. Add 2 mL of 85% H3PO4 and titrate immediately with the KMnO4 solution from the burette. Add KMnO4 drop wise until a pink color persists for about 30 seconds. Record the final burette reading in the space provided for sample #1. Repeat the procedure for the other two samples. CALCULATIONS: The number of moles of KMnO4 used in the titration is equal to the product of the molarity of the KMnO4 solution and its volume in liters used in the titration. From the balanced equation, one mole of KMnO4 is chemically equivalent to five moles of iron. EXAMPLE CALCULATION

Suppose 28.43 mL of 0.0155 M KMnO4 were required to titrate the Fe2+ in 0.998g of an iron sample. moles KMnO4 Moles of KMnO4 = 0.02843 liter x 0.0155 = 0.000441 mole 1 liter Moles of Fe2+ = 0.000441 mole KMnO4 x

5 moles Fe2+ 1 mole KMnO4

= 0.00220 mole Fe2+

55.85 g Fe2+ Grams of

Fe2+

= 0.123 g Fe2+

= 0.00220 mole Fe x 1 mole

0.123 g

Fe2+

Fe2+

% Fe2+ =

x 100% = 12.3% 0.998 g Sample

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CHEM-102

Name _________________________

Experiment #3

Date __________________________

DATA: STANDARDIZATION OF KMnO4 % Fe2+ of the known sample: _________________

#1

#2

#3

Mass of known iron sample

_______

_______

_______

Mass of Fe2+ in the sample

_______

_______

_______

Initial buret reading

_______

_______

_______

Final buret reading

_______

_______

_______

Volume KMnO4 used

_______

_______

_______

Molarity of KMnO4

_______

_______

_______

Average Molarity KMnO4

______________

Molarity KMnO4 Standard

___0.015M_____

% Error KMnO4

______________

CALCULATIONS: (Hint: Reverse the calculation for finding percent iron)

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PERCENT IRON IN UNKNOWN SAMPLE Calculated Average Molarity of KMnO4 (from page 5): ________________ Unknown #: ________________ #1

#2

Mass of unknown sample

_______

_______

_______

Initial buret reading

_______

_______

_______

Final buret reading

_______

_______

_______

Volume KMnO4 solution used

_______

_______

_______

Moles KMnO4 required

_______

_______

_______

Moles Fe2+ present

_______

_______

_______

Mass of Fe2+ present in sample

_______

________

_______

Percent Fe2+ in sample

_______

_______

_______

Average percent Fe2+

______________

Standard value (% Fe2+)

______________

% Error

______________

#3

CALCULATIONS:

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