Chemiluminescence of Metallic Sodium
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The following article was originally published in the journal for educators Chemia w Szkole (eng. Chemistry in School) (1/2014):
Introduction
Chemistry, as a scientific discipline, studies both the properties of substances that make up our world and the transformations they undergo. Chemical reactions can be understood as processes in which certain substances are transformed into others, manifesting changes in physical and chemical properties — such as state of matter, density, odor, or color. It’s important to note, however, that these changes are not always perceptible to our senses. This explains the common perception that such reactions merely transform matter. Yet we know that chemical processes are also accompanied by energy changes. Often, these are relatively subtle and, perhaps for that reason, are sometimes treated in education as something secondary or less important.
Chemiluminescent reactions belong to the class of exergonic reactions, in which visible light is generated at the expense of chemical energy. Exploring this phenomenon can be an engaging way to enrich chemistry education. Chemiluminescence is distinct from incandescence, which involves light emitted due to high temperature.
One might wonder why we usually limit our demonstrations to just the thermal effects of a reaction. One factor could be that chemiluminescent reactions seem less familiar. In reality, many such reactions have been thoroughly documented. A more significant issue is that the necessary reactants are often expensive, highly toxic, not easily accessible, or may require complex synthesis in a school laboratory that is typically not well-equipped. However, this is not a sufficient excuse — chemiluminescence can be demonstrated using relatively common substances.
Chemiluminescence was first observed in 1669 by the German alchemist Hennig Brand. He obtained the white allotrope of phosphorus and noted that it emitted a greenish-yellow glow, clearly visible in the dark. We now know that this luminescence results from the slow oxidation of phosphorus. Nevertheless, this reaction is not one you would want to perform yourself, as white phosphorus is extremely toxic — a lethal dose for an adult, whether ingested or inhaled, is around 0.1 g (0.0035 oz). There is also a fire hazard: finely divided white phosphorus ignites immediately, even at room temperature.
A much lesser-known example of chemiluminescence is the glow observed during the oxidation of alkali metals, primarily sodium and potassium. Sodium is commonly available in school chemistry labs, where it is used to demonstrate the reactivity of Group 1 elements. While sodium is highly reactive, it is not considered toxic under standard laboratory conditions, and its residues can be disposed of without excessive difficulty. All of these characteristics make the slow oxidation of sodium an excellent example for illustrating chemiluminescence, even though its glow is dimmer than that of phosphorus.
Experiment and Observations
Setting up the experiment is straightforward. You’ll need a small block of sodium — just a few cubic centimeters (e.g., around 0.06–0.18 in3) in volume. Sodium must always be stored under oil or kerosene. In this experiment, it’s more convenient to use sodium stored in kerosene, because the sample must be thoroughly dried using a paper towel or filter paper. Due to its reactivity, the surface of the metal is usually covered with oxides and other reaction products (Fig. 1).
Sodium is very soft, so you can slice through it with a knife to reveal the pure, silvery surface of the metal (Fig. 2).
The surface of the sodium quickly becomes dull due to its reaction with atmospheric oxygen. This process should be observed in a darkened room. After allowing your eyes a few minutes to adjust, you can see a faint but distinctly visible yellowish glow. It can be photographed (Fig. 3) with a long exposure time (ISO 400, 60 s, contrast enhanced).
We can therefore observe sodium’s chemiluminescence under fairly ordinary conditions. Note that the glow appears only on freshly exposed metal surfaces, while areas heavily coated with oxides remain dark.
You may wonder whether the effect can be intensified. As is generally known, the rate of most chemical reactions increases with temperature — and the same applies here. Higher temperatures should yield a brighter glow. To achieve this, cut a thin slice of sodium, place it on a metal sheet, and heat it on a hot plate (Fig. 4A) to about 40–50°C (104–122°F). Oxidation proceeds noticeably faster, and chemiluminescence becomes visible after only a few seconds in the dark. The effect can be seen in Fig. 4B (ISO 400, 20 s, contrast enhanced).
The intensity of the glow is uneven — it is particularly bright in the spot where the oxide layer was scraped off just before the photo was taken (Fig. 4B, marked with an asterisk), revealing fresh metal. Blowing dry air onto the sample intensifies the luminescence by supplying additional oxygen. However, do not use pure oxygen, as it may cause a sodium fire that is very difficult to extinguish!
After the experiment, any remaining metal should be returned to a container filled with oil or kerosene so it can be reused in future experiments. Smaller fragments can be safely destroyed by reacting them with water. This should be done cautiously and in small portions, due to the vigorous exothermic reaction.
Explanation
According to the law of conservation of energy, energy cannot simply appear from nothing or vanish without a trace. If the total internal energy of the products is lower than that of the reactants, the excess energy must be released into the environment. This usually takes the form of heat, but in the case of chemiluminescence, a portion of the energy is emitted as electromagnetic radiation with wavelengths corresponding to visible light. Thus, under the right conditions, white phosphorus, luminol, lucigenin, and even metallic sodium can emit visible light when oxidized.
Chemiluminescent reactions generally follow the scheme below:
The reactant(s) X transform into an intermediate [Y]* in an excited, high-energy state. This state is unstable and spontaneously converts into a lower-energy product Y. The surplus energy is released in the form of radiant energy, hν.
Research suggests that in sodium’s case, chemiluminescence most likely results from the reaction of atomic oxygen (O) with the sodium oxide radical NaO•. As this highly reactive radical oxidizes to stable sodium oxide Na2O — the final product — energy is released. Because the quantum yield of this reaction is relatively low (even by chemiluminescent standards), most of the energy is dissipated as heat. Only a very small fraction emerges as faint yellow light. In humid air, luminescence is significantly weaker, due to a competing reaction between sodium and water vapor, which produces sodium hydroxide (NaOH), a non-luminescent product.
It is worth noting that the reaction described here also occurs in nature, not just in laboratory conditions. In fact, it takes place in Earth’s atmosphere, near the boundary of the thermosphere, at an altitude of about 90km (56mi), where a layer of dispersed sodium atoms exists. The resulting radiation is extremely faint but can be detected using specialized instruments such as photomultiplier tubes.
This phenomenon appears to be analogous to the chemiluminescence of white phosphorus.
Further readings:
- Bielański A., Chemia ogólna i nieorganiczna, Warszawa PWN, 1981,
- Brandl H., Versuche zur Chemolumineszenz mit Alkalimetallen, Der mathematische und naturwissenschaftliche Unterricht, 46 (1993) 3, pp. 168-172
- Kolb C.E., Elgin J.B., Gas phase chemical kinetics of sodium in the upper atmosphere, Nature, 264 (1976), pp. 488-490
- Schofield K., The possible resurrection of the Chapman mechanism for atmospheric sodium chemiluminescence and ruminations on NaO reaction dynamics, International Journal of Chemical Kinetics, Volume 25 (1993), Issue 9, pp. 719-743
All photographs and illustrations were created by the author.
Marek Ples