In my graduate school days, I synthesized a number of simple carbohydrates. One of the problems was getting the products to crystallize because carbohydrates are notorious for holding onto the water with which they have been in contact. As a result, they tend to form syrups instead of crystals. When attempting to study molecular structure using an instrumental method called nuclear magnetic resonance (NMR) spectroscopy, syrups won’t do.
The challenge was to rid the syrup of water and hope the residue would then crystallize. But how do you do that?
The first attempt usually is to place the syrup in a glass vessel called a “desiccator” that is attached to a vacuum pump and hope to draw the water away. I wasn’t successful with that.
Back in those pre-computer days, you went to the library and searched books and journals for an answer. It took a while, but I did locate a carbohydrate text that suggested placing phosphorus pentoxide in the desiccator (as an aside, I just asked ChatGPT the same question and in half a second, I got the same answer). The key is that this compound reacts with water to form phosphoric acid and pulls water out of the syrup.
Back to the lab I went, and it worked. The search that yielded phosphorus pentoxide kindled an interest in this compound.
I learned that it forms when phosphorus reacts with oxygen — in other words, when phosphorus burns. But there is an interesting nuance here. Phosphorus exists in two different forms, called “allotropes.” In “white phosphorus,” four phosphorus atoms are joined together in a tetrahedral structure resulting in discreet P4 molecules. Due to strained bonds, this allotrope is highly reactive and ignites spontaneously in air. It is also toxic. When phosphorus atoms are linked in a long chain, we have much more stable, less toxic “red phosphorus.” It can still ignite, but requires a high temperature.
A deeper dive revealed that in the 1830s, a practical application was found for white phosphorus. The “Lucifer” was a match that in its head contained sulfur as a fuel, potassium chlorate to provide oxygen, and phosphorus as the igniter. The match head was covered with glue that was rubbed away as the match was struck on any surface, exposing the phosphorus to the air, causing it to catch fire and ignite the match.
Goodbye, flintstone. Hello, “phossy jaw.”
Phossy jaw was a horrific condition that afflicted workers in match factories who were exposed to the vapours of white phosphorus. It started with severe pain in the jaw and often led to necrosis of the bone and finally terrible facial disfigurement. When it became clear that exposure to phosphorus was an occupational hazard, workers in match factories —mostly women and young girls — went on strike to demand better conditions. This was the first times workers successfully forced public attention on a chemical workplace hazard.
The match that lit cigarettes also lit one of the first fights for chemical safety. Eventually, these early phosphorus matches were replaced by the “safety match,” which was based on some clever chemistry.
In 1844, Swedish chemist Gustaf Erik Pasch, who had studied under the guidance of renowned chemist Jöns Jacob Berzelius, patented a safety match that was designed to light only when struck on a special surface. It solved the problem of Lucifers spontaneously igniting in one’s pocket.
Pasch’s solution was to split the chemistry into two parts. The match head contained sulfur as fuel and potassium chlorate as oxidizer, both bound with glue to the matchstick. The matchbox featured a striking surface composed of red phosphorus and an abrasive of glass powder and sand. Striking the match on this surface created friction that converted a tiny amount of red phosphorus into white phosphorus that immediately ignited and caused the chemicals on the match head to burst into flame.
By this time, I was deeply immersed in the chemistry of phosphorus and noted that just about every publication about this element began with the fascinating story of its discovery.
In 1669, German alchemist Hennig Brandt was searching for a method to make gold. The yellow colour of urine caught his attention, and he wondered whether it might actually be due to gold. He decided to boil a sample to drive off liquids and see what was left behind. Brandt was absolutely astounded when he was left with a residue that glowed in the dark. He thought that perhaps he had discovered the much sought after “philosopher’s stone” that turned substances into gold. He had not. But he had discovered phosphorus.
Urine contains various phosphates in which oxygen atoms are linked to phosphorous. Phosphates, which originate in the diet, are used by the body to make essential biochemicals, such as DNA. Any excess is secreted in the urine. With intense heat, some of the carbon-containing compounds in urine break down to form elemental carbon that can then extract oxygen from phosphates, forming carbon dioxide, and leave elemental phosphorus behind.
It was this residue that stunned Brandt with its glow. The glow is due to the flame as the newly formed phosphorus ignites in contact with air and is converted into phosphorus pentoxide.
In just about every text, the story of the discovery of phosphorous is accompanied by a reproduction of The Alchemist Discovering Phosphorus, a 1771 painting by Joseph Wright.
It depicts an alchemist kneeling in front of his distillation apparatus as if worshipping the glowing substance in his flask. The alchemist is not named, but by the time Wright painted the picture the story of Brandt was well known. There is no known portrait of Brandt, so the figure in the painting comes from Wright’s imagination.
And that is how I went from the problem of crystalizing my 2-deoxyglucose to the painting by Wright, a reproduction of which now hangs in my office.
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