Strange Glow: Toxic Chlorine and the Chemiluminescence of Excited Nitrogen
| Polish version is here |
The following article was originally published in the journal for educators Chemia w Szkole (eng. Chemistry in School) (4/2022):
Anyone familiar with my work knows how much I value physical and chemical processes that can illustrate and explain scientific phenomena in both engaging and visually striking ways. One such category of reactions involves chemiluminescence, which consistently captures the interest of students and educators alike, not only because of its clear educational value but also due to its beautiful and often surprising visual effects.
Chemiluminescence is the emission of electromagnetic radiation in the visible range through a mechanism other than heating a substance to a high temperature. One of my personal fascinations is the synthesis and observation of chemiluminescent reactions, as well as the study of factors that influence them and the potential for practical applications. Photo 1 presents some of the better-known chemiluminescent compounds from my private laboratory collection. These include lophine C21H16N2, lucigenin C28H22N4O6, tris(2,2'-bipyridyl)ruthenium(II) chloride [Ru(bpy)3]Cl2, and luminol C8H7N3O2 [1] [2].
All of the mentioned compounds emit light of a characteristic wavelength, which corresponds to a specific color, when oxidized in specific solvents (Photo 2).
The substances described, while relatively simple to synthesize and well within the capabilities of a university laboratory, may present greater difficulties in a school or amateur lab. This is often due to the high cost of the precursors required for their preparation. Fortunately, the phenomenon of chemiluminescence can also be demonstrated using chemicals that are much easier to obtain or produce. White phosphorus; alkali metals such as sodium and potassium; singlet oxygen 1O2; polyphenols naturally present in green tea; certain organomagnesium and organosilicon compounds (such as Wöhler’s siloxene Si6O3H6, synthesized from silica SiO2); and even potassium permanganate KMnO4, which is commonly available in pharmacies, can all be used to generate visible light through chemical reactions [3] [4] [5].
Many more chemiluminescent substances are known, including naturally occurring ones responsible for bioluminescence. In Poland, this phenomenon can be observed in several beetle Coleoptera species, such as the common glow-worm Lampyris noctiluca, the lesser glow-worm Phosphaenus hemipterus, and the fire beetle Phausis splendidula [6].
Another fascinating example of bioluminescent organisms is the marine bacterium Aliivibrio fischeri, which is entirely harmless to humans under normal conditions. I successfully isolated it in my home laboratory from the surfaces of marine animals such as fish and shrimp. The process required numerous attempts and a great deal of patience. When cultivated on a suitable nutrient medium, the bacteria formed circular colonies that emitted a vivid blue luminescence (Photo 3).
In addition to the more or less well-known chemiluminescent reactions, there are others that are rarely mentioned in popular literature. This is often due to a lack of available information in domestic sources, and sometimes even in international scientific publications. In some cases, the reason may be the reaction’s limited visual appeal or the high cost and low availability of the necessary reagents.
To help bridge this gap, I would like to encourage the Reader to explore a reaction that, despite producing a clearly visible glow (in my opinion comparable to or slightly weaker than that observed in singlet oxygen reactions), remains virtually unknown. What makes it especially appealing is the fact that the required reagents are extremely inexpensive. That said, while the reaction itself is conceptually simple, performing it safely requires careful preparation of the setup and meticulous handling of all materials involved.
Chlorine
Chlorine, symbol Cl, is a chemical element with atomic number 17 and two stable isotopes: 35Cl and 37Cl. It belongs to the halogen group and is classified as a typical nonmetal. Its name, both in Latin and in modern languages such as Polish, derives from the Classical Greek word χλωρός (chloros), meaning greenish-yellow. True to its name, chlorine is a gas with a distinct greenish-yellow color and a density more than twice that of air (Photo 4). It also has a sharp, easily recognizable odor.
In nature, chlorine occurs almost exclusively in the form of chemical compounds. It plays an important biological role as a macroelement. Chloride ions are among the main anions present in bodily fluids, and hydrochloric acid is essential for activating digestive enzymes in many animals. It is estimated that a human body weighing 70 kg (about 154 lbs) contains approximately 95 g (3.35 oz) of chlorine.
Elemental chlorine, like most gases under standard conditions, exists as diatomic molecules Cl2. In compounds, it exhibits oxidation states ranging from -I to VII.
Chlorine is highly chemically reactive, though less so than fluorine. The name "halogens" reflects this reactivity, as it comes from the Greek words ἁλός (halos, salt) and γένος (genos, to produce). It readily reacts with most other elements, forming compounds known as chlorides. At 25 °C (77 °F), one liter (~34 fl oz) of water can dissolve about 2.3 liters (~81 fl oz) of the gas, producing what is called chlorinated water, a reagent commonly used in laboratory practice [7].
Chlorine is commonly used in water treatment plants during the final stage of purification. When it reacts with water, it forms chlorous acid HClO and hydrochloric acid HCl(aq), both of which have strong disinfectant properties. However, if the chlorination process is not properly controlled, chlorine can react with residual organic matter in the water to form toxic chloroform CHCl3.
Chlorine was also tested as a chemical weapon during World War I [by the Germans, author’s note], but it was eventually replaced by agents that were more effective and, arguably, even more devastating [8].
In organic chemistry, free chlorine is commonly used as an oxidizing agent, although chlorinated water is often preferred over the gaseous form. Chlorine also serves as a widely applied substituent, capable of replacing hydrogen atoms in organic compounds. Thanks to this versatility, it plays a key role in the production of a wide range of materials, including plastics, antiseptics, dyes, insecticides, petroleum products, pharmaceuticals, textiles, solvents, and many others.
Elemental chlorine was first isolated in 1774 by Carl Wilhelm Scheele, who produced it by reacting manganese(IV) oxide MnO2 with hydrochloric acid. At the time, Scheele mistakenly believed he had created a compound containing oxygen. This was due to the prevailing belief that all acids must contain oxygen, a misconception that existed before the recognition of oxygen-free acids. As a result, the new substance was thought to be an oxide of an unknown element and was provisionally named Murium or Muriaticum [9].
It wasn’t until 1810 that Humphry Davy demonstrated that the substance isolated by Scheele was not an oxide, but an element in its own right. He gave it the Latin name chlorum [10]. Shortly after its discovery, Jędrzej Śniadecki was among the first to advocate for recognizing chlorine as a distinct chemical element. The Polish name for chlorine was later proposed by Filip Walter.
Chlorine can still be produced in the laboratory today using the same method first demonstrated by Scheele, as shown in the reaction below:
| (1) |
Another option is to use potassium permanganate, treating it with hydrochloric acid in the same way as in the previous reaction:
| (2) |
In my experiments, however, I chose a different approach by using a reaction in which an acid reacts with calcium hypochlorite, Ca(ClO)2, a compound commonly known as chloride of lime. This substance is widely used to prepare disinfectant solutions, particularly for swimming pools, and is also employed in the bleaching of paper, textiles, and other materials. The reaction is shown below:
| (3) |
This method is the most efficient in terms of the amount of chlorine produced relative to the acid used. In my experience, it also proceeds in the most controlled manner. By adjusting the acid dosage carefully, it’s easy to achieve a steady and moderate flow of chlorine gas. In Photo 5, the greenish tint of the gas above the reaction mixture is clearly visible. This is free chlorine, which is then directed through an appropriate outlet tube.
It is important to avoid using rubber or certain plastic tubing, as chlorine gas can quickly degrade these materials, creating potential safety hazards. Silicone tubing is generally a good choice, but it should always be carefully checked for leaks before each use.
In any experiment involving chlorine, it is crucial to understand that this element is a powerful irritant to both the respiratory system and mucous membranes. Inhalation can cause pulmonary edema and, at sufficiently high concentrations, may even be fatal. Its sharp, distinctive odor becomes detectable in air at concentrations as low as 3.5 ppm. Exposure to levels above 800 ppm is considered lethal [11]. The permissible long-term exposure limit, averaged over an eight-hour workday, is 0.7 mg/m3, while the maximum allowable short-term concentration is 1.5 mg/m3. Because of these risks, strict safety protocols are absolutely essential. All experiments involving chlorine gas must be conducted under a fully functional fume hood.
Chlorine’s high reactivity, which is responsible for much of its hazardous nature, can be demonstrated through a particularly spectacular experiment. To perform it, you’ll need fresh flowers, with the most vibrant colors producing the most dramatic effect. In my trials, I used blossoms picked directly from my garden. The first set of tests focused on the inflorescences, or flower heads, of the rough oxeye, Heliopsis helianthoides. Each flower was suspended using a piece of insulated copper wire and placed inside a glass container (Photo 6).
The container was covered with a large glass Petri dish and then slowly filled with chlorine gas. The flower remained in this atmosphere for about 15 minutes. Afterward, it was carefully removed and exposed to fresh air to allow any residual chlorine to safely dissipate. The results of this experiment are shown in Photo 7.
As seen in the case of the oxeye inflorescence, the outer florets, which resemble petals and serve a similar visual function, are naturally yellow (Photo 7A). After exposure to chlorine gas, they became almost completely decolorized (Photo 7B). The oxeye’s inflorescence has a relatively fleshy structure, so the decolorization was incomplete. It’s likely that chlorine did not have enough time to penetrate the deeper tissues during the experiment, which is why a faint yellow hue remains visible. When using the more delicate flowers of Hosta species, the results are even more striking. The naturally purplish-pink petals (Photo 8A) turned nearly colorless after just three minutes in a chlorine-rich environment (Photo 8B). The experiment also revealed a noticeable weakening of the plant tissue’s mechanical strength following exposure to the gas.
Chlorine had an almost destructive impact on the inflorescences of garden phlox, Phlox paniculata. This species comes in many color varieties; in my experiment, I used purple-flowered specimens (Photo 9A).
In this case, visible discoloration of the petals occurred within just a few dozen seconds, accompanied by almost complete wilting and loss of the flower’s natural structure (Photo 9B).
This experiment clearly demonstrates that chlorine is a highly reactive gas, capable of rapidly oxidizing and breaking down pigments such as anthocyanins and other compounds present in plant tissues. Generally, the more water the tissue contains, the faster the bleaching effect occurs, as chlorine is significantly more aggressive in moist conditions than in dry ones.
Of course, chlorine affects not only pigments but also reacts with a wide range of organic and inorganic substances. This is why chlorine-based bleaches can easily damage fabrics.
Besides its striking visual impact, this demonstration of chlorine’s properties serves an important educational purpose. After witnessing such an experiment, few would doubt the need for strict caution when handling this substance. With that understanding, we can now turn our attention to the main topic: chemiluminescence involving chlorine gas.
Cold Light
For this experiment, chlorine will be generated using reaction (3). To carry it out, you will need the following reagents:
- calcium hypochlorite Ca(ClO)2,
- hydrochloric acid HCl(aq) (approx. 18% w/w),
- concentrated ammonium hydroxide NH3(aq) (30% w/w),
- sodium hydroxide NaOH [12].
As you can see, nearly all the substances used in this experiment carry potential hazards. Calcium hypochlorite is a powerful oxidizer that can become explosive when mixed with strong reducing agents. It also releases toxic chlorine gas upon contact with water, acids, or bases. Both hydrochloric acid and sodium hydroxide solutions are highly corrosive and can cause severe skin and eye damage. Ammonium hydroxide is a strong base as well, and the ammonia it emits is a respiratory irritant and toxic at high concentrations. Due to these risks, strict safety precautions are essential. All preparation and the experiment itself must be conducted under a properly functioning fume hood, with suitable personal protective equipment in use.
Before starting the procedure, you need to build the experimental setup. A schematic of the apparatus is shown in Fig. 1.
The apparatus consists of several interconnected components. Separatory funnel a contains 50 cm3 (approximately 1.7 fluid ounces) of hydrochloric acid at about 18% concentration. It is mounted in a rubber stopper, sealed and insulated with paraffin, which caps the suction flask b holding 15 g (approximately 0.53 ounces) of calcium hypochlorite. Together, these form the chlorine generator. The generated gas flows through silicone tubing c into the longer inlet of a Dreschel bottle d filled with concentrated ammonia solution. The outlet of this vessel leads into a second Dreschel bottle e, containing either 30% sodium hydroxide solution or, preferably, wood shavings soaked in that solution. Finally, the gas is safely vented into the fume hood through another silicone tube.
The assembled setup is shown in Photo 10. If you wish to prepare the system ahead of time, it's advisable to temporarily seal the tube connecting the generator to the first Dreschel bottle, for example with a clamp. This connection can then be opened just before or shortly after initiating chlorine production. In the latter case, be especially careful to prevent excessive pressure buildup. Regardless of timing, it is essential to ensure that all tubes, glassware, and connections are both airtight and unobstructed, as leaks or blockages could cause dangerous overpressure and potentially damage the apparatus.
It’s best to add the ammonia solution to the Dreschel bottle just before starting the experiment (Photo 11).
At this point, it should be clear that the main site of the reaction, the one where the key process occurs, is the Dreschel bottle containing the aqueous ammonia solution. The second bottle, filled with sodium hydroxide, serves to absorb most of the unreacted chlorine. However, because small amounts of chlorine gas can still escape, and given its high toxicity, the outlet must be directed safely into a fume hood.
To begin the experiment, first darken the room. Then, gently open the stopcock on the separatory funnel to allow the acid to react with the hypochlorite. Once the addition is complete, close the stopcock and observe the generation of chlorine gas. After a short delay, the gas begins to bubble rapidly through the ammonia solution in the first Dreschel bottle. During this process, brief flashes of yellow light become visible (Photo 12).
The flashes of light continue as long as chlorine flows vigorously and there is still a sufficient concentration of ammonia present. Simultaneously, gas bubbles can be seen throughout the entire volume of liquid in the Dreschel bottle.
The experiment should not be prolonged, as extended exposure may lead to the formation of nitrogen trichloride NCl3, a hazardous and potentially explosive compound. Once chemiluminescence has been observed, the chlorine generator should be promptly disconnected from the ammonia vessel, and any remaining gas should be directed into the scrubber. After the reaction has stopped and chlorine production ceases, allow the entire apparatus to ventilate under the fume hood before proceeding with cleanup.
Photo 13 shows the same experiment, this time performed in an Erlenmeyer flask.
Explanation
Interestingly, this reaction was once commonly demonstrated in school and university laboratories as a classic example of ammonia oxidation by chlorine to produce free nitrogen. Albrecht and his team noted, somewhat surprisingly, and I share their astonishment, that although the reaction was widely used for educational purposes, the distinctive chemiluminescence accompanying it was rarely acknowledged by either instructors or students [13].
In an alkaline environment, the reaction is believed to proceed according to:
| (4) |
| (5) |
As shown, the first reaction involves ammonia reacting directly with chlorine, while the second proceeds via the hypochlorite anion [14]. When performed under controlled conditions, the reaction does not produce significant amounts of nitrogen trichloride (NCl3).
The exact mechanism behind the yellow chemiluminescence observed in this reaction remains unclear. Some researchers suggest it may involve the excitation of chloramine (ClNH2), a species that can form under these conditions. However, the emission maximum associated with this pathway is around 690 nm, which corresponds to red light [15]. This clearly does not match the yellow glow observed during the experiment.
There is, however, growing evidence that the luminescence is linked to the excitation of nitrogen molecules formed in the reaction [16]. The nitrogen produced is believed to enter a highly excited triplet state (T2). The yellow light (hν) likely results from an allowed transition from this excited triplet state (T2) to a lower triplet ground state (T1), followed by a non-radiative transition to the singlet ground state (S0):
| (6) |
| (7) |
The mechanism described above is fundamentally similar to chemiluminescence involving singlet oxygen, with one important distinction. In the case of oxygen, the ground state is actually a triplet form (3O2) in which the molecule contains two unpaired electrons, making it a radical. The excited form, by contrast, has all its electrons paired and exists in a singlet state (1O2). There are two distinct singlet states, which differ in the way electrons are distributed within the π*2p molecular orbitals. It’s important to note that oxygen is a notable exception: its most stable form is the triplet ground state, while the singlet states, with fully paired electrons, are actually higher in energy and correspond to excited configurations.
Additionally, it’s worth noting that the following reaction occurs in the chlorine scrubber:
| (8) |
I encourage everyone to try these experiments for themselves, provided that all proper safety precautions are strictly followed.
References:
- [1] Ples M., Synteza i chemiluminescencja lofiny - zimne światło, muzyka i migdały (eng. Synthesis and chemiluminescence of lophine - cold light, music, and almonds), Chemia w Szkole (eng. Chemistry in School), 5 (2020), Agencja AS Józef Szewczyk, pp. 44-47 back
- [2] Ples M., Pulsujące światło - chemiluminescencyjne oscylacje (eng. Pulsating Light - Chemiluminescent Oscillations), Chemia w Szkole (eng. Chemistry in School), 3 (2019), Agencja AS Józef Szewczyk, pp. 44-48 back
- [3] Ples M., Chemiluminescencja metalicznego sodu (eng. Chemiluminescence of Metallic Sodium), Chemia w Szkole (eng. Chemistry in School), 1 (2014), Wydawnictwo EduPress, pp. 5-7 back
- [4] Ples M., Całkiem niezwykła herbatka (eng. A Rather Unusual Tea), Chemia w Szkole (eng. Chemistry in School), 4 (2015), Agencja AS Józef Szewczyk, pp. 6-9 back
- [5] Ples M., Fiolet świeci - chemiluminescencja powszechnie dostępnego związku manganu (eng. Glowing Purple – Chemiluminescence of a Common Manganese Compound), Chemia w Szkole (eng. Chemistry in School), 6 (2018), Agencja AS Józef Szewczyk, 16-19 back
- [6] Ples M., Iskrzyk - żywa latarnia (eng. Firefly - A Living Lantern), Biologia w Szkole (eng. Biology in School), 4 (2021), Forum Media Polska Sp. z o.o., pp. 52-56 back
- [7] Bielański A., Podstawy chemii nieorganicznej, Wydawnictwo Naukowe PWN, Warszawa, 2002, p. 563 back
- [8] Ellison D.H., Handbook of Chemical and Biological Warfare Agents - Second Edition, CRC Press, 2007, pp. 567-570 back
- [9] Weeks M. E., The discovery of the elements. XVII. The halogen family, Journal of Chemical Education. 1932, 9 (11): pp. 1915-1939 back
- [10] Eichstaedt I., Księga pierwiastków, Wiedza Powszechna, Warszawa, 1973, p. 180 back
- [11] Sitarek K., Chlor - dokumentacja dopuszczalnych wielkości narażenia zawodowego, w: Podstawy i Metody Oceny Środowiska Pracy, 55 (1/2008), pp. 73-95, 2008 back
- [12] Seidl M., Chemolumineszenz mit Ammoniak und Chlor, w serwisie: https://www.chem-page.de, dostępne online: https://www.chem-page.de/experimente/chemolumineszenz-mit-ammoniak-und-chlor.html [dostęp 01.08.2022] back
- [13] Albrecht S., Brandl H., Zimmermann T., Anorganische Chemilumineszenz. Traditionelle Experimente in neuem Licht, Chemie in unserer Zeit, 2008, 42 (6), pp. 394-400 back
- [14] Bray W. C., Dowell C. T., The reactions between chlorine and ammonia, Journal of the American Chemical Society, 1917, 39(5), 905-913 back
- [15] Maeda Y., Takenaka N., Chemiluminescence determination of trace amounts of ammonia and halogen species in the environment, International Society for Optics and Photonics: Bellingham (Washington, USA), 1993, pp. 185-193 back
- [16] Brown R., Winkler C. A., The Chemical Behavior of Active Nitrogen, 1970, 9(3), pp. 181-196 back
All photographs and illustrations were created by the author.
Marek Ples