074 Potentiometric Titration

Expt 074 -- Potentiometric Titration

Titration is a quantitative measuring procedure in which a liquid solution is added to a mixture until some distinctive feature signals an endpoint. The titration of Fe²⁺ with KMnO₄ is studied. As the reaction proceeds, the ratio of Fe³⁺/Fe²⁺ increases and the electrical potential changes. After the endpoint, the ratio of MnO₄^-/Mn²⁺ determines the potential. The potential of an electrode placed in this reaction mixture is compared to the potential of a Cu/Cu²⁺ reference system.

Introduction

Suppose an electrochemical cell is created by setting up an experiment according to the figure below.
This cell corresponds to a chemical reaction that can be written as
2 Fe³⁺ + Cu --> 2 Fe²⁺ + Cu²⁺
The voltage read by the voltmeter is a function primarily of the difference in the standard electrode potentials and secondarily of the ratio of the concentrations. What if the ratio of concentrations is changed, say by oxidizing some of the Fe²⁺ to Fe³⁺ with KMnO₄? In this case, the meter reading will change. Measuring changes in these cell potentials is the basis of potentiometric titrations. Once essentially all of the Fe²⁺ is reacted, the cell potential is determined by the MnO₄^-/Mn²⁺ couple versus copper.

Safety

The chemicals used are toxic. KMnO₄ and H₂SO₄ are corrosive. Wear goggles and apron. Avoid ingesting the chemicals. Avoid contact with the chemicals. Wash spills with water. Wash hands after the experiment.

Procedure

Cut all but 3 cm of the stem from a standard plastic transfer pipet.

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Place a small number of pieces of dried polyacrylamide gel on to a piece of paper. Use an unbent paper clip to insert a small piece into the end of the tube, and push it about 1.5 cm into the tube.

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Fill a well elsewhere in the plate conveniently removed from the reaction well with 0.1 M Na₂SO₄. Slowly squeeze the bulb, insert it into the Na₂SO₄ solution, and then release slowly as liquid rises in the plastic tube into the bulb covering the small piece. Once the gel swells, motion of the fluid in the plastic tube is impaired.

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Use a razor to cut a small hole (5-8 mm diameter) in the top of the bulb. The hole must be smaller than the diameter of your Cu coil.(see below)

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Use a transfer pipet to removes excess Na₂SO₄.

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Take a 12-15 cm length of solid copper wire and form a coil around a pen or pencil.

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Fill the bulb 2/3rds full with 0.5 M CuSO₄. This apparatus is the reference cell. Put a small amount of water in one well. Between titrations, store this reference cell with the tip down in the water.

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Clamp one end of a multimeter to the copper wire protruding from the bulb of the copper sulfate/copper reference cell.

Place a piece of platinum wire in one well of a 24-well plate. [Nichrome wire is a substitute; it reacts at the end, however. Also, it must be pretreated.] Bend the wire so that it fits snugly in the center of the cell. Clamp the other lead of the multimeter to the wire.

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Both cells are connected to the multimeter now.

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Switch the meter to a voltage scale. Turn it on. The reading should be 000 mv, or just a few mv (neglect sign). Trapped air bubbles defeat the purpose of the bridge. Place the tip of the bulb of reference cell into the reaction cell. The meter should now give a reading.

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Add 20 drops of 0.2 M Fe(NH₄)₂(SO₄)₂ to the platinum electrode well (the reaction well). Add 5 drops of 3 M H₂SO₄ to the well with the platinum electrode.

Stir using the reference cell. Note and record the meter reading. (It may be unstable at this point.)

Add 1 drop of 0.050 M KMnO₄ to the reaction well. Stir using the reference cell. When the reading becomes stable, record the reading. Note and record any evidence for reaction.

Continue the process of adding one drop of KMnO₄, stirring, waiting for a (fairly) stable reading, and recording. Note and record any evidence for reaction.

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Continue adding KMnO₄ dropwise, with stirring, until the cell turns pinkish. Then add dropwise 5 drops of KMnO₄. After each drop, stir, record the voltage, and record observations.

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Disconnect the meter. Remove the reference cell. Remove the electrode from the reaction well and wash the electrode with a stream of distilled water.

Discard the contents of the reaction wells into a disposal beaker provided by the instructor (Use a transfer pipet to suck out the contents from the well and transfer them to the disposal beaker or jar. Then rinse the entire plate at the sink.)

Questions

Write a balanced chemical equation of the reaction of Fe²⁺ and MnO₄^-.
The cell potential changes more at the beginning and the end of the titration than during the middle. Explain this observation.
Explain the relationship between the color change and the jump in cell potential.
A colorless substance was titrated with another colorless substance. Determine the endpoint from the voltage data. The voltage after the addition of each drop is: 1drop...268,2...278,3...284,4...288,5...296,6...302,7...305,8...310,9...321,10...329,11...340,12...370,13...556,14...625,15...677
Fe²⁺ solutions air oxidize easily. Calculate the actual concentration of Fe²⁺ in the solution titrated.

Handout Makeup

Name ___________________________ Class _______

Teacher __________________________

SmallScale 074 Potentiometric Titration

Watch the movies.

H1. State the purpose of the gel in this experiment.

Use this data to answer the questions.

0 59
1 -278
2 -268
3 -278
4 -284
5 -288
6 -295
7 -296
8 -302
9 -307
10 -310
11 -321
12 -320
13 -329
14 -340
15 -370
16 -548 Solution turns dark pink.
17 -619
18 -666
19 -694
20 -723
21 -737
22 -750
23 -759
24 -762

Curriculum-

This is an advanced topic. It is very useful in AP chemistry, or in a second year chemistry class. It is very helpful when teaching the Nernst equation.

Safety-

Time-

Teacher Preparation: 20 minutes

Class Time: 40 minutes

Materials-

5 cm platinum wire
- or
Nichrome wire and 6 M HNO₃
- (If nichrome is used, the instructor should immerse the wires in 6 M HNO₃ for two minutes, and then immerse the treated nichrome wires in water. This treatment is necessary before each titration. The titration changes the coating on the surface of the nichrome wire.)
15 cm copper wire
1 mL of 3 M H₂SO₄ -- (Use the stock solution.)
2 mL of 0.2 M Fe(NH₄)₂(SO₄)₂ -- (Add 7.84 g Fe(NH₄)₂(SO₄)₂•6H₂O to 20 mL distilled water. Add 2 drops 1 M H₂SO₄. Add enough distilled water to bring the total volume to 100 mL.)
3 mL of 0.050 M KMnO₄ -- (Add 0.79 g KMnO₄ to enough water to make 100 mL of solution. Stir often until dissolved.)
2 mL of 0.5 M CuSO₄ -- (Add 12.5 g CuSO₄¥5H₂O to 70 mL distilled water. Stir until dissolved. Add enough distilled water to bring the total volume to 100 mL.)
1 mL of 0.1 M Na₂SO₄ -- (Dissolve 1.420 g of Na₂SO₄, sodium sulfate in water. Dilute to 100 mL with water.)
standard plastic transfer pipet (for use in constructing a reference cell)
scissors to cut the pipet stem and make a hole in the top of the bulb
pencil (or pen, for making coil)
SoilMoist (pieces of dried polyacrylamide gel)
toothpicks
voltmeter (high internal impedance; for example, Micronata 22-169, LCD Digital Auto/Manual-Ranging Pocket Multimeter from Radio Shack)
alligator clips or other devices to attach meter probes to electrodes

Disposal-

Discard the salt bridges with ordinary trash. Collect the copper solutions separate from the reaction mixtures. The copper solutions may be discarded at the sink with large volumes of water. Treat the reaction mixtures with 3% H₂O₂. Filter the MnO₂ produced, and discard this with ordinary solid waste.

Lab Hints-

This is a "fussy" experiment. Be certain that a closed electrical circuit exists. The salt bridge is particularly likely to present trouble. Check for air bubbles. Displace air bubbles using a transfer pipet with a very thin stem. Fill that pipet with 0.1 M Na₂SO₄. Squeeze the pipet to remove air from the stem; insert the stem below the trapped air bubble and squeeze the bulb while removing the pipet. This will displace air bubbles.
Although paper and thread may be used as salt bridges, they cause trouble. In particular, the KMnO₄ tends to react with them.
As the titration proceeds, small amounts of the titrant may become trapped in pockets in the cell near the bridge or the electrodes. These pockets cause the cell potential to drift.
If Nichrome is used, once the Fe²⁺ is used up the electrode itself becomes susceptible to attack. If nichrome is used, the instructor should immerse the wire in 6 M HNO₃ for a minute or two, and then immerse the treated electrode in water. This treatment is necessary before each titration. The titration changes the coating on the surface of the nichrome wire.
After the reaction is over, added KMnO₄ brings about another reaction:

2 MnO₄^- + 3 Mn²⁺ + 2 H₂O --> 5 MnO₂ + 4 H⁺
When this reaction takes place, the reaction cell turns brown and murky.
The change at the endpoint is very dramatic -- several hundred millivolts.

Data Table-

0 59
1 -278
2 -268
3 -278
4 -284
5 -288
6 -295
7 -296
8 -302
9 -307
10 -310
11 -321
12 -320
13 -329
14 -340
15 -370
16 -548
17 -619
18 -666
19 -694
20 -723
21 -737
22 -750
23 -759
24 -762

Answers-

Q1. Write a balanced chemical equation for the reaction of Fe²⁺ and MnO₄^-. A1. Half-reactions: 5 ( Fe²⁺ --> Fe³⁺ + e^- ) MnO₄^- + 8 H⁺ + 5 e^- --> Mn²⁺ + 4 H₂O Complete balanced reaction: MnO₄^- + 8 H⁺ + 5 Fe²⁺ --> Mn²⁺ + 4 H₂O + Fe³⁺
Q2. The cell potential changes more at the beginning and the end of the titration than during the middle. Explain this observation. A2. At the beginning and end of the titration, the ratio [Fe²⁺]/[Fe³⁺] is changing most dramatically. It begins with a number much larger than unity (say 100), and ends much smaller (say 0.01). In the middle of the titration, the relative amounts of change are smaller.
Q3. Explain the relationship between the color change and the jump in cell potential. A3. After the endpoint, the cell potential is determined by the [Mn²⁺]/[MnO₄^-] ratio and the corresponding cell potential. The literature values are: Cu²⁺ + 2 e^- --> Cu 0.337 v Fe³⁺ + e^- --> Fe²⁺ 0.771 v MnO₄^- + 8 H⁺ + 5 e^- --> Mn²⁺ + 4 H₂O 1.51 v
Q4. A colorless substance was titrated with another colorless substance. Determine the endpoint from the voltage data. The voltage after the addition of each drop is: 1drop...268,2...278,3...284,4...288,5...296,6...302,7...305,8...310,9...321,10...329,11...340,12...370,13...556,14...625,15...677 A4. 13 drops....566 volts (Pick the point where the voltage changes rapidly.
Q5. Fe²⁺ solutions air oxidize easily. Calculate the actual concentration of Fe²⁺ in the solution titrated. A5. 20 drops Fe²⁺ solution reacts with 16 drops of 0.050M KMnO₄ in one set of sample data. (16 drops KMnO₄/20 drops Fe²⁺) x (.050 moles KMnO₄/liter) x (5 moles Fe/1 mole KMnO₄) = 0.20 M Fe²⁺ or 16 drops*0.05 M * 5 KMnO₄ = 1 Fe²⁺ *20 drops * UnknownM (16*0.05*5)/20 = 0.20 M Fe²⁺

Makeup Ans.-

H1. State the purpose of the gel in this experiment. A1. The gel is a salt bridge. The gel separates the copper reference cell from the titration reaction. Electrical contact is maintained, but only the concentrated nitrate solution is transported between the reference electrode and the reaction well.

Literature Data-

Standard electrode potentials

Cu²⁺ + 2 e^- --> Cu 0.337 V

Fe³⁺ + e^- --> Fe²⁺ 0.771 V

MnO₄^- + 8 H⁺ + 5 e^- --> Mn²⁺ + 4 H₂O 1.51 V

Key Words 1-

electrochemical cell, electrode potential, Nernst equation, salt bridge, electrode, reference cell, titration

Elements-

Fe Mn Cu Pt

Cu²⁺ + 2 e^- --> Cu	0.337 V
Fe³⁺ + e^- --> Fe²⁺	0.771 V
MnO₄^- + 8 H⁺ + 5 e^- --> Mn²⁺ + 4 H₂O	1.51 V