Let There Be Light!

Introduction

Turn on the light inside students’ minds with this colorful series of chemiluminescence demonstrations!

1. Cool Light—Demonstrates basic chemiluminescence using luminol. Two solutions are simultaneously poured into a large funnel. As the two solutions mix they undergo a chemical reaction and chemiluminescence begins.
2. Energetic Light—Demonstrates chemiluminescence as well. However, compare the longevity of this reaction to that of cool light.
3. Fluorescent Dye—Set up four beakers on a light box. When exposed to white light the solutions appear one color but in the presence of a black light they appear another.
4. Flame Test—Students observe light emitted by different metal salts as the electrons in metal ions move from excited states back to the ground state.

Concepts

• Chemiluminescence
• Oxidation–reduction
• Catalyst
• Fluorescence
• Absorbance
• Transmittance
• Emission
• Atomic emission
• Excited vs. ground states
• Wavelength and energy of light
• Flame tests

Experiment Overview

Cool Light
A wire glows, a candle burns—the production of light and heat are common to many chemical reactions. But when light is produced without heat, that’s cool! Actually, it’s called “cool light,” and the oxidation of luminol using hydrogen peroxide provides a classic demonstration of this amazing phenomenon.

Energetic Light
Two solutions are combined to produce a beautiful chemiluminescence demonstration that lasts approximately 10 minutes.

Fluorescent Dyes
Color is a result of the interaction of light with matter. The color that a solution appears to the human eye can change depending on the nature of the light source used to illuminate it. In this demonstration, four solutions that appear one color under visible light will “change” colors when exposed to an ultraviolet (black) light.

Flame Tests
Just as a fingerprint is unique to each person, the color of light emitted by an element heated in a flame is also unique to each element. In this experiment the characteristic color of light emitted by calcium, copper and sodium ions will be observed.

Materials

Safety Precautions

Hydrogen peroxide is an oxidizer and skin and eye irritant. Sodium hydroxide solution is corrosive, very dangerous to eyes, and skin burns are possible. If heated to decomposition or in contact with concentrated acids, potassium ferricyanide may evolve poisonous hydrogen cyanide fumes. Energetic Light solution contains sodium and is a corrosive liquid. Potassium ferricyanide solution is a mild irritant. Contact with strong acids may liberate toxic hydrogen cyanide gas; avoid contact with strong acids. Ethyl alcohol is flammable and a dangerous fire risk. Addition of denaturant makes ethyl alcohol poisonous—it cannot be made nonpoisonous. Dye solutions will easily stain hands and clothing; avoid all contact with skin and clothing. Do not look directly at the black light; its high-energy output can be damaging to eyes. Copper(II) chloride is highly toxic by ingestion; avoid contact with eyes, skin and mucous membranes. Fully extinguish the wooden splints by immersing them in a beaker of water before discarding them in the trash to avoid trash can fires. Wear chemical splash goggles, chemical-resistant gloves and a chemical-resistant apron. Wash hands thoroughly with soap and water before leaving the laboratory. Please review current Safety Data Sheets for additional safety, handling and disposal information.

Disposal

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 solutions from Cool Light may be disposed of down the drain with an excess of water according to Flinn Suggested Disposal Method #26b. The waste solution from Energetic Lightmay be flushed down the drain with copious amounts of water according to Flinn Suggested Disposal Method #26b. The Fluorescent Dyes solutions may be disposed of down the drain according to Flinn Suggested Disposal Method #26b.

Prelab Preparation

Cool Light

1. Prepare solution A by adding 0.1 g of luminol and 50 mL of 5% sodium hydroxide solution to approximately 800 mL of distilled or deionized (DI) water. Stir to dissolve the luminol. Once dissolved, dilute this solution to a final volume of 1000 mL with DI water.
2. Prepare solution B by adding 0.7 g of potassium ferricyanide and 15 mL of 3% hydrogen peroxide to approximately 800 mL of DI water. Stir to dissolve the potassium ferricyanide. Once dissolved, dilute this solution to a final volume of 1000 mL with DI water.
3. Set up the demonstration equipment as shown in Figure 1.
   {12770_Preparation_Figure_1}

Energetic Light

1. Add 70 mL of the 0.6% potassium ferricyanide and 7 mL of 3% hydrogen peroxide solution to one 250–mL beaker.
2. Add 70 mL of energetic light solution to a second 250-mL beaker. Note: The energetic light solution is very viscous.

Fluorescent Dyes

1. Obtain four 600-mL beakers.
2. Add 500 mL of tonic water into one of the 600-mL beakers. This will be known as beaker 1.
3. To a second 600-mL beaker (beaker 2), add 15 mL of the 1% fluorescein solution. Dilute the fluorescein solution by adding enough distilled or deionized water to reach the 500-mL mark on the beaker.
4. Add 5 mL of the 1% eosin Y solution to the third 600-mL beaker. Dilute the eosin Y solution by adding enough ethyl alcohol to reach the 500-mL mark on the beaker. Stir. Note: Eosin Y is soluble in water but the fluorescence is not nearly as strong in water as it is in ethyl alcohol.
5. Add 1 mL of the 1% rhodamine B solution to the fourth 600-mL beaker. (Use a graduated cylinder or add 15 drops from the dropping bottle.) Dilute the rhodamine B solution by adding enough distilled or deionized water to reach the 500-mL mark on the beaker. Stir.
6. View the four solutions under classroom lights and note the color.
7. Darken the room as much as possible and turn on the ultraviolet light source. Hold the black light behind the beakers and view the colors of the emitted light.

Procedure

Cool Light

1. Turn down the lights. The room should be as dark as possible.
2. Pour solution A and Solution B into the large funnel simultaneously. As the two solutions mix, chemiluminescence begins.
3. As the reaction progresses, it can be enhanced by adding small amounts of potassium ferricyanide and 5–10 mL of 5% sodium hydroxide solution into the flask.

Energetic Light

1. Darken the room completely.
2. While stirring with a stirring rod, add the potassium ferricyanide/hydrogen peroxide solution to the energetic light solution.
3. Observe the chemiluminescence. Stir the solution occasionally to prolong the light reaction.

Flame Tests

1. Fill a 250-mL beaker about half-full with distilled or deionized water. Soak three wooden splints in the beaker containing distilled or deionized water several hours or overnight.
2. Fill a second 250-mL beaker about half-full with tap water. Label this beaker “rinse water.”
3. Label three weighing dishes Ca, Cu and Na. Place one scoopful (about 0.5 g) of each solid metal chloride into the corresponding weighing dish.
4. Light the laboratory burner.
5. Dip the soaked end of one of the wooden splints in one of the metal chlorides, then place it in the flame. Observe the color of the flame. Allow the splint to burn until the color fades. Try not to allow any of the solid to fall into the barrel of the laboratory burner. If necessary, repeat the test with the same splint and additional salt.
6. Immerse the wooden splint in the “rinse water” to extinguish it, then discard it in the trash.
7. Instruct students to record their observations for the flame color produced by the metal chloride in the data table on the Let There Be Light Worksheet.
8. Repeat steps 5–7 for the other two metal chlorides. Record your observations for the flame color produced by each metal ion in the data table.

Student Worksheet PDF

12770_Student1.pdf

Teacher Tips

• This kit contains enough chemicals to perform the Cool Light demonstration seven times: 1 g of luminol, 5 g of potassium ferricyanide, 500 mL of 5% sodium hydroxide and 105 mL of 3% hydrogen peroxide.
• This demonstration is especially appealing if it is set up so the students can see the mixture through spiraling clear plastic tubing. This type of apparatus gives a large surface area for light to be emitted as well as providing a flowing effect along with the luminescence—increasing the overall visual impact.
• Use only distilled or deionized water when preparing the solutions. Hard water and softened water contain high concentrations of ions (such as chloride ions) that may interfere with the excited state of the luminol and prevent chemiluminescence.
• This kit contains enough chemicals to perform the Energetic Light demonstration seven times: 500 mL of energetic light solution, 50 mL of 3% hydrogen peroxide solution and 500 mL of potassium ferricyanide solution.
• Make sure you stir the energetic light solution while you are adding the 0.6% potassium ferricyanide solution. This will add to the longevity of the chemiluminescence.
• Chemiluminescence is not a bright light; it is more of a glow. The darker the room, the brighter the glow will appear.
• This kit contains enough chemicals to perform the Fluorescent Dyesdemonstration seven times: 1 L of tonic water, 30 mL of Rhodamine B solution, 30 mL of Eosin Y solution, and 500 mL of 95% ethyl alcohol. The tonic water may be reused in subsequent demonstrations.
• The difference in the colors of the solutions under the classroom lights compared to the black light is most obvious in a completely darkened room. Try to extinguish all light sources.
• The solutions will last all day, even a whole week; however, some evaporation will occur. The solutions will keep for an extended period of time if the beakers are covered with Parafilm^®.
• The tonic water does not have to be carbonated for the fluorescence to occur—it will still fluoresce if it is flat.
• The prescribed dilutions listed above are not strict. The same effects are noticeable over a range of solution concentrations; however, the fluorescence is not as easily observed in dilute solutions. For best results, solutions more concentrated than those suggested above are not recommended.
• All students or other observers should wear protective eyewear. Caution them not to look directly at the black light.
• This kit contains enough chemical to perform the Flame Tests demonstration at least seven times: 50 g of calcium chloride, 20 g of copper(II) chloride, 50 g of sodium chloride and 30 wooden splints.
• Provide students with the information found on Data Table 2 in thte Fluorescent Dyes demonstration in order to answer the questions on the Let There Be Light! Worksheet.
• Other metal ions that give bright, characteristic flame test colors include lithium, strontium, barium and potassium.

Sample Data

Flame Tests

Results Table

Answers to Questions

Cool Light

1. Describe what happened in this demonstration.

   Two solutions were poured through a large funnel simultaneously. The mixture of the two solutions began to glow.

2. Oxidation is necessary for luminol to luminesce. The chemicals used in this experiment were 5% sodium hydroxide, 3% hydrogen peroxide, and potassium ferricyanide. Which of these do you think served as an oxidizing agent?

   Hydrogen peroxide is the oxidizing agent in this demonstration.

3. In chemiluminescence a molecule is produced in an “excited” state (i.e., the electrons are at a high energy level). The electrons in the molecule then return to their stable state (i.e., lower energy level.) Explain how this is linked to the production of light.

   When an electron drops to a lower energy level, energy must be released. This energy is released in the form of light.

4. Define chemiluminescence. Give an example of chemiluminescence found in nature.

   Chemiluminescence is a process in which light is produced through chemical reactions. An example found in nature is the firefly.

Energetic Light

5. Describe what happened in this demonstration. How was the color of the mixture different than the cool light mixture?

   Two solutions were mixed together and resulted in a yellow-green mixture. The cool light mixture was blue.

6. How else were the results of the Energetic Light demonstration different than the Cool Light demonstration?

   Other than the color another difference was the duration of the chemiluminescence. Energetic light lasted much longer than cool light.

Fluorescent Dyes

7. Draw the four beakers. Label each beaker with its contents, the color of the solution under normal white light, and the color of the solution under black light.
   {12770_Answers_Figure_7}
8. The visible spectrum detected by the human eye ranges from about 400 to 700 nm. UVA light (black light) transmits in a range from about 320 to 400 nm. Explain why we cannot see the light from a black light as we can see a normal light.

   Black lights transmit higher energy light than the light that is within our visible range, therefore we cannot see that particular kind of light. But often black lights will transmit wavelengths in the low 400s. These wavelengths appear violet to the human eye, giving black light its purple glow.

9. Fluorescence occurs when a substance absorbs a photon from a light source. The energy from that photon causes an electron to move to an “excited” state (higher energy level). As that electron returns to its ground state it releases another photon with a particular wavelength. Explain how this relates to the “colorful glow” seen when a substance fluoresces.

   The glow is caused by the energy that is released by the electron relaxing from a high energy level to a low energy level. If the photon that is released at this time has a wavelength that is within the visible spectrum, then we can see the colorful glow it causes.

Flame Tests

10. Using the data provided by your instructor, record the approximate wavelength of light emitted for each metal ion in the results table.

    See above data table.

11. Convert each wavelength in the results table from nanometers to meters. Show one sample calculation in the space below and record all values in the results table.

    Sample calculation for calcium:
    λ in meters = 450 nm x (1 m/1 x 10⁹ nm)
    λ in meters = 4.5 x 10^–7 m

12. Using the equation ΔE = hc/λ, calculate the average energy corresponding to the observed flame color for each metal. Show one sample calculation in the space below and record all values in joules in the results table.

    Sample calculation for calcium

    {12770_Answers_Equation_2}
13. A glass rod was heated in a burner flame and gave off a bright yellow flame. What metal ion predominates in the glass rod?

    The yellow flame is characteristic of sodium ions, which is a key component in the composition of glass.

Discussion

Cool Light

Chemiluminescence is defined as the production or emission of light that accompanies a chemical reaction. Light emission results from the conversion of chemical energy into light energy due to changes in the composition of a chemiluminescent material. The “flame test” colors observed when different metal salts are burned in a Bunsen burner flame are examples of a type of chemiluminescence called pyroluminescence. The glow of solid phosphorous in air is another classic example of chemiluminescence—light, along with some heat, is produced when the phosphorus undergoes an oxidation reaction. The oxidation of luminol (3-aminophthalhydrazide) in this demonstration illustrates a type of “cool light” chemiluminescence in which little or no heat is produced.

The light-producing chemical reactions of luminol were discovered by H. O. Albrecht in 1928. Since that time numerous procedures have been developed to produce light using luminol. Experiments have shown that the following “ingredients” are necessary for luminol to exhibit chemiluminescence—a basic (alkaline) pH, an oxidizing agent, and a catalyst. In this demonstration, the oxidizing agent is hydrogen peroxide, the catalyst is the iron(III) cation in potassium ferricyanide, and the sodium hydroxide is used to maintain the basic pH required for the reaction to occur.

Oxidation of luminol and the resulting chemiluminescence occurs in the following sequence of reactions:

1. Sodium hydroxide acts as a base and converts luminol (structure I) into a dianion.
2. Hydrogen peroxide oxidizes the dianion form of the luminol to the aminophthalate ion (structure II), which is produced in an excited electronic state.
3. The excited aminophthalate ion decays to a lower energy ground state and gives off light in the process. The emitted light has a wavelength of 425 nm, which is the blue region of the visible spectrum.
   {12770_Discussion_Figure_2}

Energetic Light

Chemical luminescence, or chemiluminescence, is a process by which chemical energy is converted into light energy. Both Cool Light and Energetic Light solutions display chemiluminescence. However, chemiluminescence lasts much longer in Energetic Light solution due to the presence of calcein (also known as fluorexon) in this solution. Calcein is a fluorescent dye that acts as a sensitizer. As luminol is produced in an excited state light is emitted as the electrons fall back to ground state. In the presence of calcein the reaction is prolonged as the luminol transfers the energy to calcein. Calcein has the ability to self-quench which yields a large oxidation potential, therefore prolonging luminescence. In general terms, the chemical reaction generates a product in an intermediate excited state (the electrons are at a higher energy level than in the ground state). When the electrons fall from that excited state to the more stable ground state, energy is released in the form of light.

A + B → C + D* → D + light

The chemiluminescence reaction follows an energy diagram like the diagram. Fluorescent Dyes

Absorption and Transmission of Visible Light
The four solutions appear different colors under the normal classroom lights. The tonic water is colorless, the fluorescein solution is yellow-green, the eosin Y solution is yellow-orange, and the rhodamine B solution is pinkish-red. They are each composed of different molecules—molecules that absorb different wavelengths of light. In general, a green solution looks green to the human eye because it is transmitting green light. When white light is shined through this solution, the molecules in the solution absorb some of the wavelengths of the light and transmit others. All nongreen wavelengths of light will be absorbed by the green solution to some extent, although red light will be absorbed the most. The red photons hit the solution and are absorbed by the molecules in the solution. They do not make it through the solution, and hence, we do not see a red color from this solution. In contrast, green photons are not absorbed by the molecules in the green solution so they pass right through the solution, and we see a green color.

How do we know that the green solution absorbs the red wavelengths of light? Red and green are complementary colors—they are across from each other on the color wheel.

In general, colors opposite each other on the color wheel are complementary colors. For example, by looking at Figure 4, it can be seen that violet and yellow are complementary colors. Therefore, it can be assumed that a yellow solution absorbs violet light and transmits yellow light. The color wheel and the idea of complementary colors can be used as a first estimation of the wavelengths that are absorbed by a substance based on its color. The following table lists the wavelengths associated with each of the colors in the visible spectrum and their complements. The representative wavelength can be used as a benchmark for each color. For example, instead of referring to green as light in the wavelength range 500–560 nm, one could simply say that green light is 520 nm.
Fluorescence
Luminescence is the emission of radiation (light) by a substance as a result of absorption of energy from photons, charged particles, or chemical change. It is a general term that includes fluorescence, phosphorescence and chemiluminescence. Fluorescence is different from other types of luminescence in that is it restricted to phenomena in which the time interval between absorption and emission of energy is extremely short. Therefore, fluorescence occurs only in the presence of the exciting source. This is different from phosphorescence, which continues after the exciting light source has been removed. In this demonstration, the exciting source is the UV black light.

In fluorescence, when a light source is shined on a material, a photon is absorbed. The energy from the photon is transferred to an electron that makes a transition to an excited electronic state. From this excited electronic state, the electron naturally wants to relax back down to the ground state. When it relaxes back down to the ground state, it emits a photon (symbolized by the squiggly arrow in Figure 5). This relaxation may occur in a single step or in a series of steps. If it occurs in a single step, the emitted photon will be the same wavelength as the exciting photon. If the relaxation occurs in a series of steps emitting a photon along the way, the emitted photon will have a greater wavelength (lower energy) than the exciting photon. If the emitted photon’s wavelength is in the visible portion of the spectrum, we observe a colorful, glowing effect. Emission of this form is termed fluorescence. This process is practically instantaneous so the fluorescence is observed as soon as the exciting source is present, and it disappears as soon as the exciting source is removed. The fluorescent glow is brighter than the color of the solution seen under normal classroom lights because light is being emitted from the solution, not just transmitted through it.

Absorption Curves and Color
Information about the absorption and emission curves of each of the solutions in this demonstration can be inferred from the observations made during the demonstration—that is, what wavelengths of light these solutions absorb and emit.

First consider the fluorescein, eosin Y, and rhodamine B solutions.

• These solutions appear colored to the human eye under the normal classroom lights. Recall that normal classroom lights give off white light which is composed of all the visible wavelengths of light. Therefore, these solutions must absorb some wavelengths of visible light while transmitting others—the color of the solution is the transmitted color in each case. Each of these colored solutions has an absorption (and transmission) peak in the visible region of the electromagnetic spectrum (400–700 nm).
• When the normal classroom lights are turned off and the black light is shined on the solutions, they fluoresce. Under these conditions, the solutions are not being hit with visible light, but instead are being hit with UVA light (320–400 nm). In each case, when a molecule in the solution is hit with ultraviolet photons, the molecule absorbs an ultraviolet photon and promotes an electron up to an excited state. This electron then relaxes back down to the ground state in a series of steps emitting a visible photon along the way. It is evident that the photon is in the visible region of the spectrum because the fluorescence can be seen with the human eye. Therefore, the molecules in each of the solutions must have an absorption peak in the UVA portion of the electromagnetic spectrum with a corresponding emission peak in the visible portion of the spectrum.

Clearly, each of these solutions has two absorption peaks—one in the visible and another in the UVA portion of the spectrum. If the transmitted wavelength of visible light is not the same wavelength as the emitted photon during fluorescence, the solution will appear to be two different colors under the two different light sources.

Now consider the tonic water solution. Tonic water appears colorless to the human eye under the normal classroom lights. Therefore, it must not absorb any wavelengths of visible light. Consequently, in contrast to the three solutions discussed above, it does not have an absorption peak in the visible region of the spectrum. But, under the UVA black light, it is blue! When hit with ultraviolet light, one of the ingredients in tonic water, quinine, absorbs an ultraviolet photon and emits a visible photon in return. The human eye can see this visible photon, and therefore this solution appears to be colored when viewed under the black light.

It is evident from these examples that color is not an inherent quality of a substance, but instead, a result of the interaction of light with matter. If the wavelength of the light changes, the interaction, and hence the resulting color, may also change.

Flame Tests

When a substance is heated in a flame, the atoms absorb energy from the flame. This absorbed energy allows the electrons to be promoted to excited energy levels. From these excited energy levels, there is a natural tendency for the electrons to make a transition or drop back down to the ground state. What an electron makes a transition from a higher energy level to a lower energy level, a particle called a photon is emitted (see Figure 6). Both absorption and emission of energy are quantized—only certain energy levels are allowed. An electron may drop all the way back down to the ground state in a single step, emitting a photon in the process. Alternatively, an electron may drop back down to the ground state in a series of smaller steps, emitting a photon with each step. In either case, the energy of each emitted photon is equal to the difference in energy between the excited state and the state to which the electron relaxes. The energy of the emitted photon determines the color of light observed in the flame. The flame color may be described in terms of its wavelength, and Equation 1 may be used to calculate the energy of the emitted photon. ΔE is the difference in energy between the two energy levels in joules (J), h is Planck’s constant (h = 6.626 x 10^–34 J•sec), c is the speed of light (c = 2.998 x 10⁸ m/sec) and λ (lambda) is the wavelength of light in meters. The wavelengths of visible light are given in units of nanometers (1 m = 1 x 10⁹ nm) (see Table 1 in Fluorescent Dyes).

The color of light observed when a substance is heated in a flame varies from one substance to another. Because each element has a different spacing of electron energy levels, the possible electron transitions for a given substance are unique. Therefore, the difference in energy between energy levels, the exact energy emitted photon, and the corresponding wavelength and color are unique to each substance. As a result, the colors observed when a substance is heated in a flame may be used as a means of identification.

The Visible Portion of the Electromagnetic Spectrum
Visible light is a form of electromagnetic radiation. Other familiar forms of electromagnetic radiation include γ-rays, X-rays, ultraviolet (UV) radiation, infrared (IR) radiation, microwave radiation and radio waves. Together, all forms of electromagnetic radiation make up the electromagnetic spectrum. The visible portion of the electromagnetic spectrum is the only portion that can be detected by the human eye—all other forms of electromagnetic radiation are invisible.

The visible spectrum spans the wavelength region from about 400 to 700 nm. Light of 400 nm is seen as violet and light of 700 nm is seen in red. According to Equation 1, wavelength is inversely proportional to energy. Therefore, violet light is higher energy light than red light. As the color of light changes, so does the amount of energy it possesses. Table 1 lists the wavelengths associated with each of the colors in the visible spectrum. The representative wavelengths may be used as a benchmark for each color. For example, instead of referring to green light in the wavelength range of 500–560 nm, we may approximate the wavelength of green light at 520 nm. An infinite number of shades of each color may be observed.

References

Cesa, I. Flinn ChemTopic™ Labs, Volume 6, Atomic and Electron Structures, 2004, p 89.

Harvey, E. N., A History of Luminescence. The American Philosophical Society: Philadelphia, PA, 1957; p 5.

Huntress, E. H.; Stanley, L. N.; Parker, A. S., J. Chem. Educ., 1934, 11, 145.

Shakhashiri, B. Z. Chemical Demonstrations: A Handbook for Teachers of Chemistry, Vol. 1: The University of Wisconsin: Madison, 1983, pp 161–167.

Shakhashiri, B. Z. Chemical Demonstrations: A Handbook for Teachers of Chemistry, Vol. 1; University of Wisconsin: Madison, 1985; p 189.