The Photochemical Reduction of Thionine

Description: An acidified solution of thionine and iron(II) sulfate is exposed to a source of intense light. The color of the solution turns from purple to colorless within a few seconds. The purple color is restored when the light source is removed. By adding various reactants to this complex, redox equilibrium, the rates at which the color changes occur can be modified.

Materials:


Hazards: Handle 6 M sulfuric acid, concentrated nitric acid and 6 M sodium hydroxide with caution.

Procedure: Part I:

  1. The following should be done prior to the demonstration. In a 1-liter or 2-liter beaker, thoroughly mix together 10 mL of 0.001 M thionine solution, 10 mL of 6 M sulfuric acid, sufficient distilled water to bring the volume to 500 mL, and finally about 2.0 g of iron (II) sulfate heptahydrate. Prepare a second solution in a similar way in another beaker, but do not add the iron (II) sulfate. Put 100 mL of this second solution into a 250 mL beaker for later use in Part II of the procedure.
  2. Place the first beaker (containing the iron (II) sulfate) on the stage of an overhead projector. Turn on the lamp and ask your students to observe what happens. The solution becomes colorless within a few seconds. Turn off the lamp and observe that the color is restored rather quickly (several seconds).
  3. Repeat the cycle a few times, noting the time it takes for both the bleaching process and the restoration of the purple color. Which reaction, the bleaching or restoration of color, occurs at a faster rate?
  4. Place the second beaker on the overhead projector, and turn on the lamp. Note that the purple color does not fade. Add about 2.0 g of iron (II) sulfate heptahydrate to this beaker while it is still on the overhead and stir. The purple color will now disappear.
  5. To illustrate that the bleaching reaction in the first beaker is initiated by light, and not by heat, place several thicknesses of aluminum foil across half of the overhead projector and under the beaker so that only about half of the solution is exposed to the bright light. Only that half of the solution which is exposed to the light will turn colorless. (A portion of the solution may also be heated on a hot plate to show that heating to the boiling point of the solution does not cause bleaching.) Discuss the reaction in terms of bond breaking/forming reactions, oxidation-reduction, equilibrium, and rates.

Discussion: In acid solution, iron(II) reduces thionine in the presence of intense light. This is a strikingly visible example of the conversion of light to chemical energy, and the demonstration could lead to a discussion of photosynthesis. The reaction is reversible, and the rates of both the forward (light) and reverse (dark) reactions are easily measured.

The equilibrium reaction can be represented by the equation:
[THIONINE REACTION]

In the oxidized form of the dye, there is a conjugation of double bonds across the three rings. The electronic excitation energy for this molecule is within the visible region of the spectrum. In the reduced form, the double bond on the nitrogen in the middle ring is no longer present, and conjugation is restricted to the two outer rings. Any electronic transitions (such as pi bonding to pi antibonding) now have energies in the ultraviolet region. The use of a control solution containing only the thionine and sulfuric acid shows that the bleaching process requires the presence of iron(II).

Procedure: Part II:

  1. Pour 100 mL of the solution into each of four 250-mL beakers. Use one of these beakers as a reference. Place the second beaker next to the reference beaker on the overhead projector, and turn on the lamp. When the solutions are colorless, turn off the lamp. Quickly add 5 mL of 0.1 M iron (III) nitrate to one of the beakers. Note that the color is restored more quickly to purple in this beaker.
  2. Place the third beaker next to the reference beaker on the overhead projector, and turn on the lamp. When the solutions are colorless, turn off the lamp. Quickly add 5 mL of 1 M sodium dihydrogen phosphate to one of the beakers. Note that the purple color is not restored as quickly in this beaker; also the original color intensity is not restored.
  3. To the fourth beaker add concentrated nitric acid (several mLs) until the solution turns a light red; upon standing, the solution may slowly turn colorless unless too much nitric acid has been added. The red color indicates a third form of the thionine which has been oxidized.
  4. Use the 100 mL of the solution containing only thionine and sulfuric acid in a fifth 250-mL beaker, and add 10 mL of 6 M NaOH. The color changes from purple to red. If the NaOH is added to a beaker that contains iron (II), a precipitate of iron hydroxide will initially form, but eventually it will turn red if enough NaOH is added.

Discussion: Addition of iron(III) after the bleaching reaction increases the rate of restoration of the purple color by shifting the equilibrium to the left. Iron(III) forms stable complexes with phosphate ions. The phosphate complex of Fe(III) has a lower reduction potential than the uncomplexed Fe(III); formation of the complex shifts the above equilibrium to the right. The addition of a strong base produces the deprotonated form of thionine, which has a red color. To the surprise of the students, the addition of an oxidizing nitric acid does not become colorless, but initially shifts the equilibrium to the left producing this same red form. As the reaction develops, the solution may eventually return to the purple, then colorless form. The type of reduction which thionine undergoes - gain of H+ ions and electrons - is characteristic of several biochemical reduction reactions, such as the transfer of electrons from nicotinamide adenine dinucleotide (NADH) to molecular oxygen in oxidative phosphorylation.

Disposal: Mix the solutions together and adjust the pH to be within a range of 5-9 by adding sodium bicarbonate or citric acid, as necessary. Solutions can then be flushed down the drain.

Reference: Lawrence J. Heidt, "The Photochemical Reduction of Thionine," J. Chem. Ed., 26, 525 (1949).

Bob Cairo, Woodrow Wilson Summer Institute, Princeton, 1988. Modifications by Mark Case, CHEM 6 Team TORCH Binder, 1995