Phosphorus, the Light of Alchemy
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The following article was originally published in the journal for educators Chemia w Szkole (eng. Chemistry in School) (2/2019):
Alchemy was an ancient practice that wove together elements now belonging to several modern scientific disciplines, chiefly chemistry and physics, but also art, psychology, and others. The most widely recognized goals of alchemists included discovering a method to transmute lead and other base metals into gold, creating a panacea capable of curing all diseases, and preparing an elixir of immortality. At the center of these quests stood the legendary philosopher’s stone, a mysterious substance whose elusive pursuit consumed the lifetimes of many.
Today, the field is often ridiculed for the naïveté (from a modern perspective) of many of its practices and claims. Alchemists did not, as a rule, employ anything resembling what we now recognize as the scientific method. Yet it must be acknowledged that the work of some was systematic and did yield tangible results. Indeed, the groundwork laid by alchemy paved the way for the emergence of true science in the modern sense, including chemistry itself. A similar transition occurred in astronomy, which gradually distinguished itself from the mystical framework of astrology.
One of the notable achievements of alchemists was the discovery of numerous new substances. Among them was the element phosphorus P, first isolated in 1669 by the German Hennig (or Henning) Brand [1]. The English painter Joseph Wright, working in the eighteenth century, immortalized this discovery in one of his canvases (Photo.1).
In this work, the artist presented his vision of Brand’s laboratory. Notice the retort containing the newly isolated phosphorus, glowing with an otherworldly light. This detail was not Wright’s invention, although he did exaggerate the luminescence considerably.
We now know that the glow is caused by the slow oxidation of the newly discovered element. That is why it was named phosphorus (from the Greek "phosphoros," meaning “light bearer”). Interestingly, had the Latin root been chosen instead, the substance might have been called lucifer, carrying precisely the same meaning (from Latin "lux" - "light" and "ferre" - "to bring"). But let us set etymology aside for now.
Would today’s practitioner of the chemical arts, leaving alchemy to the historians, wish to explore at least some of the intriguing properties of phosphorus? The answer, I believe, can only be yes, especially since doing so does not require sophisticated laboratory facilities. Still, it must be emphasized with absolute clarity that any experiments with this substance demand the utmost caution, given its dangerous properties.
Different Forms
Phosphorus is a classic example of allotropy, a phenomenon in which a single element can exist in different modifications, each with distinct physical and chemical properties. These allotropes may vary in their crystalline structures or in the number of atoms per molecule [2].
Transformations between allotropes are first-order phase transitions, involving a discontinuous change in a state function such as entropy, while thermodynamic equilibrium is maintained. Such processes drive the system toward minimizing its free energy. An element may therefore exist in two allotropic forms at the same temperature, although under specific conditions one form is more stable than the other.
Allotropy is observed in many elements. Carbon is perhaps the best known, occurring as graphite, diamond, nanotubes, and several other forms.
Phosphorus occurs in four allotropes:
- white phosphorus (sometimes referred to as yellow),
- red phosphorus (amorphous),
- violet phosphorus,
- black phosphorus.
For our purposes, the white and red forms are the most relevant. The others are chemically far less reactive. Violet phosphorus can be obtained either by heating red phosphorus in a vacuum above 500°C (932°F) or by crystallizing white phosphorus from molten lead, whereas black phosphorus is prepared by heating white phosphorus in the absence of oxygen at 220°C (428°F) under a pressure of 12,000 atmospheres [3].
Red and White
White phosphorus is the most chemically reactive allotrope. It appears as a whitish or yellowish wax-like solid with a melting point of 44°C (111°F) and a density of 1.8 g/cm3 [4]. Each molecule consists of four atoms arranged in a regular tetrahedral structure.
Although white phosphorus does not react with water, it must be carefully shielded from air because it oxidizes violently upon exposure. For this reason, it is stored under distilled water in sealed containers.
White phosphorus is extremely toxic. The lethal dose for an adult, whether ingested or inhaled, is about 0.1 g (0.0035 oz). This fact alone requires extraordinary caution in handling, ideally with heat-resistant trays made of ceramic or a comparable material. Work with both white and red phosphorus must always be carried out under strict personal and fire safety precautions.
The chemiluminescence of white phosphorus is easy to demonstrate. A fragment only a few millimeters in size can be blotted dry and placed on an inert, nonflammable surface such as ceramic or heat-resistant glass. Even when kept under water, the surface of a sample often develops an oxide film over time (Photo.2A), but this does not interfere with the observation.
In a darkened room, the sample emits a distinct glow (Photo.2B), often showing a subtle whitish-green hue.
Such demonstrations must be brief. As phosphorus oxidizes, it gradually warms and may ignite spontaneously. Solid pieces can ignite at only a few tens of degrees Celsius, and some sources report ignition occurring at temperatures as low as 20°C (68°F). In powdered form, phosphorus ignites immediately. Burning white phosphorus reaches temperatures of about 1000°C (1832°F), resists extinguishing with water, and produces highly corrosive fumes of P4O10 and its hydrolysis products. For this reason, once the glow has been observed, the sample must be promptly returned beneath the surface of water.
A historical forensic test for detecting phosphorus, including its presence in human tissues suspected of poisoning, relied on observing the chemiluminescence of vaporized white phosphorus. This method became known as the Mitscherlich test.
One of the darker chapters in the history of chemistry is the use of white phosphorus in warfare, most notoriously in incendiary bombs.
The second allotropic form of interest is red phosphorus, which occurs as a dark red powder (Photo.3) and is insoluble in all common solvents.
Unlike white phosphorus, the red allotrope is not inherently toxic. Nevertheless, samples must still be handled with care, as they may contain trace but hazardous amounts of white phosphorus as an impurity. Red phosphorus is far less reactive, does not oxidize under normal conditions, and does not need to be stored under water. Its ignition requires much higher temperatures.
Although handling red phosphorus is safer, it still requires proper precautions. A simple test demonstrates this. Place a tiny fraction of a gram on a ceramic plate and ignite it (Photo.4A). It burns less violently than white phosphorus but produces a similar white smoke of phosphorus oxides. Covering the sample with a glass vessel cuts off oxygen and extinguishes the flame (Photo.4B). Strikingly, once uncovered, the material may reignite spontaneously. Heat from combustion converts some of the red allotrope into white, which ignites readily and transfers the flame back to the red form (Photo.4C). Fires involving red phosphorus can therefore be just as difficult to extinguish as those involving white phosphorus.
White Phosphorus from Matches
White phosphorus is not readily available today, and for good reason, above all because of its extreme toxicity.
So what is an experimenter to do if they wish to experience, at least in part, what Brand might have felt in his alchemical workshop upon encountering this remarkable substance? Fortunately, there is a simple method using ordinary matches (Photo.5).
Some readers may be surprised, since it has long been known that matches with heads containing white phosphorus have not been manufactured since the late nineteenth and early twentieth centuries. Public concern, fueled by frequent poisonings and accidental fires because such matches were prone to spontaneous ignition, led to the signing of an international convention banning white phosphorus in match production (Bern, September 26, 1906). Poland acceded to this treaty in 1921 [5].
Modern matches therefore contain no white phosphorus. Instead, they rely on red phosphorus, located not in the match head but in the striking surface, the abrasive strip. This strip also contains other ingredients, such as powdered glass, which provides the necessary roughness.
To convert the red phosphorus in the striking surface into white phosphorus and observe its chemiluminescence, cut away a portion of the strip from a matchbox and carefully separate the thinnest possible layer of coated paper (Photo.6). This minimizes the unnecessary bulk of cardboard material.
Next, fold the paper into a trough with the coated side facing inward and place it on a cold metal or glass surface, such as the bottom of an inverted Petri dish (Photo.7A). Light the paper (Photo.7B) and let it to char (Photo.7C).
After the paper has burned away, a dark, resinous deposit forms on the cold glass surface (Photo.8).
This deposit contains, along with combustion products and binding agents, trace amounts of white phosphorus formed from the red allotrope in the strike surface. The quantity is minute but sufficient: in a darkened room, once the eyes have adjusted, a faint greenish glow becomes visible (Photo.9). As the outer layer oxidizes, the luminescence fades. Scraping the deposit exposes fresh phosphorus, causing the glow to brighten again. Eventually, however, all of the phosphorus oxidizes and the process comes to an end.
This method is relatively safe, since it generates only trace amounts of the hazardous allotrope. Even so, caution remains absolutely essential.
Slightly larger quantities of white phosphorus can be obtained by placing a few milligrams of the red allotrope at the bottom of a narrow test tube and displacing the air with carbon dioxide to protect the product from oxidation. The tube is then stoppered with a plug of glass wool. Gentle heating of the base of the tube with a burner converts the red allotrope into white, which vaporizes. If the walls are cooled at the same time, for example by wrapping them with wet filter paper, the vapors condense into a whitish deposit [6]. After cooling and removing the stopper, the deposit glows clearly upon contact with air (Photo.10). The apparent color differences in the photographs result from variations in camera white balance; in reality, the glow was similar to that observed in the earlier demonstrations.
It must be emphasized once again: white phosphorus is extremely dangerous! Even the smallest fragment discarded in a waste container can spontaneously ignite after some time, causing a fire. Its toxicity poses an additional risk. All residues must be rendered harmless, for example, by incineration under controlled conditions or by prolonged shaking with a solution of copper(II) sulfate (CuSO4), which converts the phosphorus into stable phosphides [7].
Explanation
The light emitted by white phosphorus arises from its oxidation, as shown by the fact that luminescence ceases under anaerobic conditions. Most chemiluminescent reactions can be described by the general scheme:
Here, substrate X is transformed into an intermediate [Y]* in a high-energy excited state. Because such states are inherently unstable, the intermediate relaxes to a lower-energy final product Y. The excess energy is released as radiant energy, hν.
For phosphorus, the element is first oxidized to lower oxides, which are then further oxidized to phosphorus(V) oxide, P2O5, more accurately represented as P4O10. This reaction accounts for the visible glow [8]. The observation that luminescence weakens or even disappears when phosphorus is exposed to pure oxygen supports this explanation: under such conditions, oxidation proceeds directly to the final oxide without generating excited intermediates.
The phenomenon resembles the chemiluminescence I have previously described during the oxidation of sodium Na [9]. The comparison is striking because the substances belong to entirely different categories: white phosphorus is a typical nonmetal (with only black phosphorus showing semiconducting behavior), while sodium is a prototypical metal.
We can therefore conclude that phosphorus displays highly unusual properties, and that chemiluminescence is neither as rare nor as exotic as it may seem. Indeed, it can be demonstrated with materials readily available in the home.
References:
- [1] Weeks M. E., The discovery of the elements (Supplementary note on the discovery of phosphorus), Journal of Chemical Education, 10 (5), 1933, p. 302 back
- [2] Gajewski W. (red.), Encyklopedia techniki - Chemia, Wydawnictwa Naukowo-Techniczne, Warszawa, 1965 back
- [3] Bielański A., Podstawy chemii nieorganicznej, PWN, Warszawa, 2002, p. 644 back
- [4] Averbuch-Pouchot M.T., Durif A., Topics in Phosphate Chemistry, World Scientific, 1996 back
- [5] Oświadczenie Rządowe o przystąpieniu Polski do konwencji międzynarodowej o zakazie używania białego (żółtego) fosforu przy wyrobie zapałek, podpisanej w Bernie dnia 26 września 1906 roku, dostępne online: http://prawo.sejm.gov.pl/isap.nsf/DocDetails.xsp?id=WDU19220190159 [dostęp: 4.03.2019] back
- [6] Ples. M., Chemiluminescencja fosforu, w serwisie: http://www.weirdscience.eu/, dostępne online: http://weirdscience.eu/Chemiluminescencja%20fosforu.html [dostęp: 4.03.2019] back
- [7] Roesky H.W., Mockel K., Doświadczenia chemiczne, Wydawnictwo Adamantan, Warszawa, 2001, p. 43 back
- [8] Pluciński T., Doświadczenia chemiczne, Wydawnictwo Adamantan, Warszawa, 1997, pp. 14-15 back
- [9] 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
All photographs and illustrations were created by the author.
Addendum
The video below shows the chemiluminescence of phosphorus:
Simple as it is, the effect is definitely worth recommending.
Marek Ples