Olympic Rings Chemistry Demonstration

Produce the colors of the Olympic rings in this easy chemistry color change demonstration. The colors of the Olympic Rings represent the colors of the flags of all the countries that participated in the 1912 Olympic Games. The Olympic Rings are blue, black, red, yellow, and green on a white background.

Perform a chemistry demonstration in which you fill empty glasses, arranged like the Olympic Rings, with “water”. The colors of the Olympic Rings appear upon adding a clear liquid to the glasses.

This demonstration illustrates coordination chemistry, precipitation reactions, and metal–ligand complex formation involving Fe³⁺. Each ring color results from a distinct chemical interaction between iron(III) and a specific ligand or counterion.

Key Takeaways: Olympic Rings Chemistry Demonstration

• The Olympic Rings demonstration showcases coordination chemistry using iron(III).
• Different ligands produce distinct colored complexes.
• Some colors form by complexation, others by precipitation.
• The demo illustrates selective chemical reactivity.
• It is suitable for general chemistry, AP Chemistry, and inorganic chemistry.

Olympic Ring Materials

You need a container of 5 grams ferric ammonium sulfate in 500 milliliters of water. In separate glasses, dissolve about half a gram of the following solids in a few milliliters of water:

• red – potassium thiocyanate
• white – barium chloride
• blue – potassium ferrocyanide
• black – tannic acid
• green – tartaric acid
• yellow (amber) – sodium hydrogen sulfite

Perform the Olympic Rings Chem Demo

Arrange the glasses to form the order of the Olympic Rings. Try to use as small a volume as possible to dissolve the solids so the glasses will appear empty. Pour the ferric ammonium sulfate into the glasses and watch the colors develop!

• red – thiocyanate ion forms a deep red complex with iron(III)
• white – barium ion reacts with the sulfate ion to form a cloudy white precipitate
• blue – ferrocyanide ion with iron(III) produces a deep blue compound
• black – tannic acid with iron(III) forms a black complex
• green – tartaric acid with iron(III) forms a greenish complex
• yellow (amber) – hydrogen sulfite ion with iron(III) produces an amber compound

Troubleshooting the Olympic Rings Demonstration

Colors are faint or slow to develop

• Increase reagent concentration slightly.
• Ensure ferric ammonium sulfate solution is fresh.
• Check that the iron solution contains Fe³⁺ and has not been reduced.

Blue ring does not form properly

• Potassium ferrocyanide may have oxidized.
• Ensure correct reagent was used, ferrocyanide and not ferricyanide.
• Verify sufficient Fe³⁺ concentration.

White ring does not appear cloudy

• Confirm sulfate ions are present in the ferric ammonium sulfate solution.
• Increase barium chloride concentration slightly.

Green or yellow turns brown

• Iron(III) may be hydrolyzing.
• Adjust solution to mildly acidic conditions to stabilize Fe³⁺.
• Use freshly prepared solutions.

Precipitate settles quickly

• Gently swirl the glass before presentation.
• Use smaller volumes to improve visual intensity.

Unexpected color mixing

• Ensure glasses are clean and free of cross-contamination.
• Use separate droppers or pipettes for each reagent.

Chemical Reactions and Balanced Equations

1. Red Ring – Iron(III)–Thiocyanate Complex

This is an equilibrium complex formation reaction.
$$\text{Fe}^{3+}(aq) + \text{SCN}^-(aq) \rightleftharpoons [\text{FeSCN}]^{2+}(aq)$$

Note:

• This is a coordination complex.
• The deep red color is due to ligand-to-metal charge transfer.
• The reaction is reversible and equilibrium-dependent.

2. White Ring – Barium Sulfate Precipitation

This is a classic precipitation reaction.
$$\text{Ba}^{2+}(aq) + \text{SO}_4^{2-}(aq) \rightarrow \text{BaSO}_4(s)$$

Brief explanation:

• Barium sulfate is highly insoluble.
• The white cloudiness results from formation of a fine precipitate.
• This is not a coordination complex, but an ionic precipitation reaction.

3. Blue Ring – Prussian Blue Formation

This produces a coordination polymer (Prussian blue).

Simplified net ionic form:
$$4\text{Fe}^{3+}(aq) + 3[\text{Fe(CN)}_6]^{4-}(aq) \rightarrow \text{Fe}_4[\text{Fe(CN)}_6]_3(s)$$

Notes:

• This forms an insoluble, intensely colored coordination compound.
• The structure is an extended lattice, not a simple 1:1 complex.
• The deep blue color is due to intervalence charge transfer between Fe²⁺ and Fe³⁺ centers.

4. Black Ring – Iron(III)–Tannate Complex

Tannic acid is a polyphenol and forms iron–tannate complexes.

This is best represented generically because tannic acid is a mixture of galloyl units:
$$\text{Fe}^{3+}(aq) + \text{tannic acid} \rightarrow \text{iron(III) tannate complex}$$

Note:

• Similar chemistry is responsible for traditional iron gall ink.
• The dark color results from strong metal–phenolate interactions.

Avoid forcing an artificial stoichiometric equation since tannic acid composition varies.

5. Green Ring – Iron(III)–Tartrate Complex

Tartrate acts as a chelating ligand.

Simplified representation:
$$\text{Fe}^{3+}(aq) + \text{C}_4\text{H}_4\text{O}_6^{2-}(aq) \rightarrow \text{Fe(tartrate)}^+(aq)$$

Notes:

• Tartrate is a bidentate ligand.
• Chelation stabilizes the complex and modifies its absorption spectrum.
• Color may vary depending on concentration and pH.

6. Yellow (Amber) Ring – Iron(III)–Hydrogen Sulfite Interaction

Hydrogen sulfite forms a coordination complex and may partially reduce Fe³⁺ under some conditions.

Simplified coordination representation:
$$\text{Fe}^{3+}(aq) + \text{HSO}_3^-(aq) \rightarrow \text{iron–sulfite complex}$$

Note:

• In strongly acidic solutions, sulfite acts as a mild reducing agent.
• Under neutral classroom conditions, the dominant process is complex formation.

Patriotic Colors Chemistry Demonstration

A related project is a patriotic color demonstration for your country’s flag. For example, in the US, use the chemicals to make red, white and blue. If you live in Brazil, prepare the chemicals to make green, yellow and blue.

Safety Considerations

This demonstration uses iron(III) salts and several reactive ligands. Although suitable for classroom use, follow proper laboratory safety procedures.

• Wear safety goggles and gloves at all times.
• Avoid ingestion or skin contact with all solutions.
• Barium chloride is toxic if ingested and should be handled with care.
• Potassium ferrocyanide is stable and does not release free cyanide under neutral conditions, but it should still be handled responsibly.
• Ferric ammonium sulfate may stain skin and clothing.
• Thiocyanate salts are harmful if swallowed.
• Dispose of all solutions according to local laboratory waste guidelines. Do not pour concentrated heavy metal solutions directly down the drain unless permitted by your institution.

If performing this demonstration with students, the instructor should prepare the solutions in advance.

References and Further Reading

• Lawrance, Geoffrey A. (2010). Introduction to Coordination Chemistry. Wiley. doi:10.1002/9780470687123. ISBN 9780470687123.
• Miessler, Gary L.; Donald Arthur Tarr (1999). Inorganic Chemistry. Prentice Hall. ISBN 978-0-13-841891-5.
• Schwarzenbach, Gerold (1957). Complexometric Titrations. Translated by Irving, Harry (1st English ed.). London: Methuen & Co.
• Zumdahl, Steven S. (2009). Chemical Principles (6th ed.). New York: Houghton Mifflin Company.