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Investigating the reaction between manganate(VII) and ethanedioate ions

Description

Use a continuous monitoring method to investigate the redox reaction between potassium manganate(VII) and ethanedioate ions.
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Potassium manganate(VII), KMnO4, is a deeply coloured purple crystalline solid. It is a powerful oxidising agent. In acidic solution, it undergoes a redox reaction with ethanedioate ions, C2O42-. The MnO4- ions are reduced to Mn2+ and the C2O42- ions are oxidised to CO2.

2MnO4-(aq) + 16H+(aq) + 5C2O42-(aq)   2Mn2+(aq) + 8H2O(l) + 10CO2(g)

In this experiment students use a continuous monitoring method to investigate how the rate of the reaction changes with the concentration of MnO4- ions. They then go on to investigate the effect of the addition of a small amount of Mn2+ on the reaction.

 


This practical activity takes around 1 hour.

burette (50 cm3)

50 cm3 volumetric flask / measuring cylinder (×5)

colorimeter with cuvettes

test tubes with rubber bungs

10 cm3 measuring cylinder

pipettes

potassium manganate(VII) solution, 0.002 mol dm-3 (no hazard)

acidified ethanedioic acid solution (CORROSIVE – causes severe skin burns and eye damage) (Note 1)

manganese(II) sulfate solution, 0.02 mol dm-3 (no hazard)

distilled water

Refer to Health & Safety and Technical notes section below for additional information.

 


Students should wear eye protection when carrying out this experiment. 

 This solution must be 0.125 mol dm-3 with respect to ethanedioic acid and 1.5 mol dm-3 with respect to sulfuric acid. 100 cm3 of this mixture can be made up by dissolving 1.125 g of ethanedioic acid (MODERATE HAZARD – harmful if swallowed, harmful in contact with the skin) in 25 cm3 of distilled water and making up to 100 cm3 with 2 mol dm-3 sulfuric acid (CORROSIVE – causes severe skin burns and eye damage).

 


 Prepare the following solutions of KMnO4 containing varying concentrations of MnO4- ions.

Using a burette, carefully transfer the required amount of the KMnO4 solution (0.002 mol dm-3) into a 50 cm3 volumetric flask or measuring cylinder and make up to 50 cm3 with distilled water.

Solution

Volume of 0.002 mol dm-3 KMnO4 solution added / cm3

Concentration of MnO4- ions in final solution / mol dm-3

1

10.0

4.0 × 10-4

2

7.5

3.0 × 10-4

3

5.0

2.0 × 10-4

4

2.5

1.0 × 10-4

5

1.0

0.40 × 10-4

 

 Place a cuvette containing distilled water into a colorimeter and using a suitable filter adjust to 0% absorbance.

 Place each of the solutions 1 to 5 into the colorimeter in turn and read off the corresponding absorbance.

 Plot a graph of absorbance (y-axis) against concentration (x-axis) - the calibration curve.

 Place 2.0 cm3 of a 0.002 mol dm-3 solution of potassium manganate(VII) in a test tube. Fit the tube with a rubber bung.

 Zero a stop clock ready for use.

 Using a small measuring cylinder, add 8.0 cm3 of the acidified ethanedioic acid solution to the test tube containing the potassium manganate(VII) solution.

Quickly stopper the tube, invert the tube to mix the contents and start the stop clock.

 Using a teat pipette, quickly transfer some of the mixture from the test tube to a cuvette and place the cuvette in the colorimeter. Measure the absorbance of the mixture in the cuvette every 20 seconds until the absorbance drops to 0.01.

 Using the calibration curve, convert the absorbance values obtained into concentrations of MnO4-.

 Plot a graph of time against concentration of MnO4- ions.

 Repeat steps e to j, but this time add 1 drop of the 0.02 mol dm-3 solution of manganese(II) sulfate to the acidified ethanedioic acid solution before mixing.

 Compare the two graphs. What effect does the addition of 1 drop of Mn2+(aq) have on the reaction?

 


The reaction between manganate(VII) ions and ethanedioate ions at room temperature is fairly slow initially but quickens as the reaction proceeds. Manganese(II) ions, Mn2+, formed as the reaction proceeds act as an autocatalyst. Once a small amount of Mn2+ ions have formed, they can react with MnO4- ions to form Mn3+ ions as an intermediate species. These then react with the C2O42- ions to reform Mn2+.

4Mn2+(aq) + MnO4-(aq) + 8H+(aq)   5Mn3+(aq) + 4H2O(l)

2Mn3+(aq) + C2O42-(aq)   2CO2(g) + 2Mn2+(aq)

In the reaction carried out in the absence of manganese(II) ions, the reaction initially proceeds slowly. Then as a few Mn2+ ions are formed, the rate increases and proceeds following first-order kinetics with respect to the manganate(VII) ions. Adding Mn2+ ions at the beginning of the reaction results in a smooth decrease in concentration with time from the outset.

As an extension task the students can go on to use their plots of concentration of MnO4- ions against time to determine the rate of the reaction for different concentrations of MnO4- ions. A plot of log(rate) against log[MnO4-] should then be a straight line with a gradient of 1 indicating that the reaction is first order with respect to manganate(VII) ions.


Page last updated October 2016

ADDITIONAL INFORMATION

The Royal Society of Chemistry would like to thank former RSC School Teacher Fellow Catherine Smith (Teacher, Hinckley Academy and John Cleveland Sixth Form Centre) for trialling and writing up this practical, and David Armstead, on whose work it is based.

Image © Shutterstock.