A colorless solution is added to each of six beakers. A second colorless solution is added and each of the six resulting solutions turns a different color of the rainbow. Add more of the first solution and the rainbow of colors disappears.
- Acids and bases
- pH indicators
Hydrochloric acid solution, 0.1 M, HCl, 1 L*
Indicator solutions: Violet, Blue, Green, Yellow, Orange, Red*
Sodium hydroxide solution, 0.1 M, NaOH, 1 L*
Beakers, 250-mL, 6
*Materials included in kit.
Hydrochloric acid solution, although dilute, is severely corrosive to eyes, skin and other tissue. Sodium hydroxide solution, although dilute, is corrosive; skin burns are possible; very dangerous to eyes. The indicator solutions contain ethyl alcohol, which is a flammable liquid and a fire risk; keep away from heat and open flame. Wear chemical splash goggles, chemical-resistant gloves and a chemical-resistant apron. Please review current Safety Data Sheets for additional safety, handling and disposal information.
Please consult your current Flinn Scientific Catalog/Reference Manual for general guidelines and specific procedures, and review all federal, state and local regulations that may apply, before proceeding. The resulting solution may be flushed down the drain with excess water according to Flinn Suggested Disposal Method #26b.
- Set up the six 250-mL beakers on an overhead projector or light box, or in front of the class.
- Add 3 drops of “violet” indicator solution to the first beaker. Add 3 drops of “blue” indicator solution to the second beaker. Continue adding three drops of each of the other indicator solutions to the appropriate beakers.
- Do not allow the solvent of each indicator solution to evaporate.
- Dilute the 0.1 M sodium hydroxide 10 to 1 to make a 0.01 M solution. Do the same with the 0.1 M hydrochloric acid.
- Add approximately 50 mL of the 0.01 M hydrochloric acid solution to each of the six beakers. All six resulting solutions should be clear.
- Add approximately 75 mL of 0.01 M sodium hydroxide solution to each beaker. Each of the six solutions should change from clear to a color of the rainbow!
- Add approximately 100 mL of the 0.01 M hydrochloric acid solution to each beaker. The solutions will once again be clear.
Precise amounts of acid and base solutions are not important. Each addition of acid or base solution must neutralize the solution in the beaker and drive the pH in the opposite direction. All solutions can be poured into the large two-liter beaker. The resulting solution will be acidic and clear.
- The indicators are dissolved in 95% ethyl alcohol. The alcohol will readily evaporate, leaving the indicator powder in the beaker—unseen to the observers of the demonstration.
- You may use drops of a more concentrated acid or base to change the solutions from colored to clear, or a more concentrated base to change the solutions from clear to colored.
Correlation to Next Generation Science Standards (NGSS)†
Science & Engineering Practices
Asking questions and defining problems
Analyzing and interpreting data
Engaging in argument from evidence
Disciplinary Core Ideas
MS-PS1.A: Structure and Properties of Matter
MS-PS1.B: Chemical Reactions
HS-PS1.A: Structure and Properties of Matter
HS-PS1.B: Chemical Reactions
Cause and effect
Scale, proportion, and quantity
Stability and change
MS-PS1-2. Analyze and interpret data on the properties of substances before and after the substances interact to determine if a chemical reaction has occurred.
HS-PS1-2. Construct and revise an explanation for the outcome of a simple chemical reaction based on the outermost electron states of atoms, trends in the periodic table, and knowledge of the patterns of chemical properties.
Answers to Questions
- Draw a diagram of the set-up. Include the chemicals that were added to the beakers. For each beaker, list the original color of the first solution, the color change after the second solution was added, and the final color after the third solution was added.
First solution: 0.01 M Hydrochloric Acid
Second solution: 0.01 M Sodium Hydroxide
Third solution: 0.01 M Hydrochloric Acid
- Given that the two chemicals added to the beakers were an acid and a base, what kind of chemical must have already been present in the beakers to produce the color changes?
An acid–base indicator must have been present in each beaker, since the solutions were different colors when a base, NaOH, was added than when an acid, HCl, was added.
- Three indicators are used in this demonstration: phenolphthalein, thymolphthalein and p-nitrophenol. Phenolphthalein is an indicator that is colorless in an acidic solution but pink-red in a basic solution. Thymolphthalein is also colorless in an acid, but blue in a base, and p-nitrophenol is colorless in an acid and yellow in a base. What indicator or combination of indicators was responsible for the color change in each beaker?
Beaker 1 – thymolphthalein and phenolphthalein
Beaker 2 – thymolphthalein
Beaker 3 – Thymolphthalein and p-nitrophenol
Beaker 4 – p-nitrophenol
Beaker 5 – phenolphthalein and p-nitrophenol
Beaker 6 – phenolphthalein
The three indicators used in this lab, phenolphthalein, thymolphthalein, and p-nitrophenol, are colorless in acidic solution. In a basic solution, phenolphthalein is red, thymolphthalein is blue and p-nitrophenol is yellow. Any color in the spectrum is possible using these primary colors.
The solvent of each indicator solution added to the beakers readily evaporates, leaving only a residue of the indicators on the bottom of each beaker. Students will not see this step of the procedure. They will only see the pouring of the acid and base solutions and the color changes. This demonstration can be done as a “magic” show. It is a great demonstration to get your students to look beyond what is happening and get them asking questions about how the disappearing rainbow occurs. Initially students might think the clear acid–base solutions contain a magical indicator. Knowing there is no such thing, have your students propose explanations for how these six colors could appear.
Shakhashiri, B. Z. Chemical Demonstrations; University of Wisconsin: Madison, WI, 1989; Vol. 3, pp. 41–46.