Chemical Education Journal (CEJ), Vol. 7, No. 2 /Registration No. 7-13/Received October 17, 2003.
URL = http://www.juen.ac.jp/scien/cssj/cejrnlE.html

e-mail: jorge.ibanez@uia.mx

INTRODUCTION
EXPERIMENTAL
RESULTS
CONCLUSIONS
ACKNOWLEDGEMENT
REFERENCES

Laboratory experiments can sometimes be used for different educational levels with only minor modifications. This paper presents a microscale redox chemistry experiment that can be used with middle school, high school, or college students depending on the depth of the concepts to be discussed and the type of experimentation selected.

INTRODUCTION

It is well known that many elements can exist in different oxidation states, and that these often have strikingly different properties. For example, Cr(VI) is very toxic while Cr(III) is an essential element. Likewise, Fe(II) is highly soluble in mildly acidic media whereas Fe(III) precipitates. S(VII) is a powerful oxidizer but S(IV) is a strong reductant. In the present experiment we study Mn in four of its oxidation states: II, IV, VI and VII. Interestingly, all of them have different color! (1). We will perform a redox titration of potassium permanganate (KMnO4) with sodium bisulfite (NaHSO3) according to the target knowledge level as follows:

a) Middle school.

Place 1 mL of the permanganate solution in each of the three small containers. Add, respectively, 1 mL of sulfuric acid, or 1 mL of water, or 1 mL of NaOH to each. Add dropwise the bisulfite solution to each beaker. Observe the change in color of the Mn(VII) solution according to the pH environment to: a) faint pink, almost transparent (in acid), b) brown (in neutral), or c) green (in base). With the observed colors and a chemistry book or handbook, find out which manganese ion was produced in each container.

b) High school.

As above, except that the experiment can be done in such a way as to promote that the students match their observations of the color of the final solution as well as the relative amount of bisulfite required, to the final oxidation state of manganese.

c) College.

As above, except that:

1. The students will be asked to balance the equations in advance
2. The experiment will be done quantitatively in the form of a titration, and
3. The students will perform a calculation to obtain the number of electrons transferred per ion of Mn(VII). With this number, they will verify the oxidation state of the manganese ions produced in each medium.

In the redox titration of potassium permanganate (KMnO4) with sodium bisulfite (NaHSO3) in different media, oxidation and reduction processes occur. The following reactions describe the process. (Physical states have been omitted for simplicity, since all species considered here are aqueous, except water, which is liquid). A species that becomes reduced gains electrons and consequently acquires a lower oxidation state. For example, in an acidic solution, Mn⁷⁺ gets reduced to Mn²⁺ by electron gain as shown in Equation 1:

MnO4- + e- + H+ --> Mn2+

(unbalanced)

(1)

Since the electrons gained by Mn⁷⁺ must have come from somewhere, for every species that gets reduced another one must simultaneously get oxidized. During oxidation, the species that gets oxidized loses electrons and consequently increases its oxidation state. For example, in this experiment the bisulfite ions are oxidized to sulfate ions (2,3):

HSO3- --> SO42- + 2e-

(unbalanced)

(2)

Since the oxidation state of sulfur in bisulfite is +4 whereas that in sulfate is +6, one concludes that each bisulfite loses two electrons.

Example: Balancing Redox Reactions in Acidic Solutions

Hydrogen ions and water are also present here, together with the species that get oxidized or reduced, and thus they can be used for balancing the corresponding half-reactions. The reaction between MnO4^- and HSO3^- in acid can be balanced by analyzing both reactions separately. Let us start with the permanganate reaction (reaction 1).

First of all, we note that the oxidation state of manganese in permanganate is 7+, whereas in the resulting manganese ion it is 2+. Thus, 5 electrons are required per every permanganate ion. To balance the oxygens one has to consider that the oxygen-containing product is water. Thus one has to add one water molecule per every oxygen in permanganate and get:

MnO4- + 5e- + H+ --> Mn2+ + 4 H2O

(3)

Now we can proceed by analyzing the hydrogens. There are 8 on the product side, and thus the same amount is required on the reactants side:

MnO4- + 5e- + 8 H+ --> Mn2+ + 4 H2O

(4)

To verify the balance, one analyzes jointly the MASS and TOTAL CHARGE balances on both sides. By doing so, we observe that the global balance is correct.

The other half-reaction (reaction 2) is easier to balance since one can note that there is one less oxygen on the bisulfite side. Thus, adding one water to the reactants and three hydrogen ions to the products, both the mass and charge of the reaction become balanced:

HSO3- + H2O --> SO42- + 2 e- + 3 H+

(5)

Note that the electrons appear on the left side in one of the half-reactions and on the right side on the other. This must always be the case since, as discussed earlier, the reduction of a species implies the oxidation of another.

To obtain the total reaction, both half reactions must be added in such a way that the total number of electrons is the same on both sides, and then we cancel them. To get to do this, the half reaction of permanganate is multiplied by 2 and that of bisulfite by 5:

2 (MnO4- + 5e- + 8 H+ --> Mn2+ + 4 H2O)

(6)

5 (HSO3- + H2O --> SO42- + 2e- + 3 H+)

(7)

At this stage we can add both reactions, the electrons get cancelled, and the total (net) balanced redox reaction is:

2 MnO4- + 5 HSO3- + 16 H+ + 5 H2O --> 2 Mn2+ + 5 SO42- + 8 H2O + 15 H+

(8)

Simplifying the waters and protons, we get:

2 MnO4- + 5 HSO3- + H+ --> 2 Mn2+ + 5 SO42- + 3 H2O

(9)

As an exercise for the higher-level students, they can now balance the corresponding reactions in neutral and basic media.

EXPERIMENTAL

Materials and Reagents

a) Middle school

1 Plastic delivery pipet
3 Small containers (e.g., beakers, Erlenmeyer flasks, etc.)
3 Stirring rods
KMnO4 (approximately 0.01 M)
NaHSO3 solution (approximately 0.02 M)
1 M H2SO4
1 M NaOH

b) High School

Same as above. A better measuring device (e.g., Microburet) is advisable.

c) College

1 Microburet (2-mL pipet)
3 10-mL Erlenmeyer flasks (or beakers)
1 Stand
1 Three-finger clamp
0.01 M KMnO4
0.02 M NaHSO3 solution
1 M H2SO4
1 M NaOH

Procedure

a) Middle school

Place 1 mL of the permanganate solution in each of the three small containers. Add drop wise the bisulfite solution to each beaker. Observe the change in color of the Mn(VII) solution according to the pH environment to: a) faint pink, almost transparent (in acid), b) brown (in neutral), or c) green (in base). With the observed colors and a Chemistry book or handbook, find out which manganese ion was produced in each container.

b) High school

Proceed as above, but count the number of drops of the bisulfite solution used in each case to reach the end-point (i.e., the clear-cut color change). For more accuracy use a better measuring device (e.g., a microburet). Divide each number of drops by the number of drops used for the titration in the basic medium. Based on these results, match the manganese oxidation state with the color observed. Verify your assignment with the aid of a Chemistry book or handbook.

c) College

With a volumetric pipet place 1.00 mL of approximately 0.01 M KMnO4 in three Erlenmeyer flasks. Add 1 mL of 1 M H2SO4 to the first one, 1 mL of deionized water to the second, and 1 mL of 1 M NaOH to the third. Fill the microburet with a 0.0200 M NaHSO3 solution and titrate each permanganate solution with it, with manual or (preferably) magnetic stirring. The purple color of the solution will suddenly become almost transparent at the end-point in the acidic medium, or brown (suspension-precipitate) in the neutral medium, or green in the basic medium. If indicated by your instructor, replicate one or more times each titration. Mn(VII) is reduced to Mn(II), Mn(IV), Mn(VI), respectively. Balance the equation for each reaction (note that the one for the acidic medium was developed above).

Write down the volume of bisulfite solution required for each titration. Calculate with the average the concentration of the permanganate solution. Then, obtain relative volumes in the acidic, neutral and basic titrations as compared to this last one. Lastly, calculate the number of electrons transferred per ion of Mn(VII). With this number, verify the oxidation state of the manganese ions produced in each medium.

RESULTS

We have performed the experiments with a microburet 10 times each, and obtained the following results (see Table 1). The concentration of the potassium permanganate solution was calculated with the average volume required in the acidic titration as follows:

1.314 mLbisulfite (2 mmolpermanganate/5 mmolbisulfite) (0.02 mmolbisulfite/mL)(1/mL)

= 0.0105 Mpermanganate

Then, the moles of electrons per mol of permanganate were calculated as follows (this calculation is shown for the acidic titration):

1.314 mLbisulfite (0.02 mmolbisulfite/mL) (2 mmole^-/1 mmolbisulfite)(1/0.0105 mmolpermanganate)

= 5.01 mol e^-/mol KMnO4

Table 1. Experimental results of the titration of KMnO4 with 0.02 M NaHSO3

	Sample	Acidic	Neutral	Basic
	1	1.35	0.90	0.25
	2	1.30	0.80	0.26
	3	1.30	0.80	0.26
	4	1.36	0.80	0.26
	5	1.30	0.90	0.26
	6	1.30	0.80	0.25
	7	1.32	0.80	0.26
	8	1.30	0.80	0.27
	9	1.31	0.90	0.25
	10	1.30	0.80	0.25
Average		1.314	0.83	0.257
Std Dev		0.022705848	0.04830459	0.00674949
Variance		0.000515556	0.00233333	4.5556E-05
Vol bisulfite/Vol bisulfite in base		5.11	3.23	1.00
mol e/mol KMnO4		5.01	3.16	0.98

It can be seen that the theoretical predictions of 5, 3 and 1 moles of electrons per mole of manganese (for the acidic, neutral and basic media, respectively) match reasonably well with the observed results.

CONCLUSIONS

Permanganate is titrated with bisulfite at different pH values to produce Mn(II), Mn(IV), or Mn(VI). We have adapted a microscale titration procedure with bisulfite that can be used as a multilevel experiment for Middle School, High School or College-level Chemistry courses.

ACKNOWLEDGEMENT

Elizabeth Garcia-Pintor and Sebastian Terrazas-Moreno assisted with the experimentation. Universidad Iberoamericana and CONACYT (Mexico) provided financial assistance. This paper is dedicated to Ronald M. Pike (retired from Merrimack College, and co-founder of the National Microscale Chemistry Center) for his long-lasting inspiration to the microscale chemistry community.

REFERENCES

1. Z. Szafran, R.M. Pike, J. C. Foster. Microscale General Chemistry Laboratory. Wiley: New York, 1993. pp. 267-275.

2. Fritz, J. S.; Schenk, G. H. Jr. Quantitative Analytical Chemistry, 2nd. Ed., Allyn and Bacon: Boston, 1969. Chapter 13.

3. Ramette, R. W. Chemical Equilibrium and Analysis. Addison-Wesley: Reading, MA, 1981. Chapter 12.

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