Phosphorus and the history of the match

In 1669, Hennig (Nicholas) Brand(t) isolated phoshorus from urine. This was one of the first non-metallic elements to have been discovered. Prior to this the only elements known included: Antimony, Bismuth, Carbon, Copper, Gold, Iron, Lead, Mercury, Silver, Sulfur and Tin. Arsenic was discovered in 1250, Phosphorus in 1669, followed by Platinum in 1735.

"The Alchemist in Search of the Philosopher's Stone" (1771)
A depiction of the discovery of the element phosphorus by German alchemist Hennig Brand in 1669.
Painted by Joseph Wright (3 September 1734 - 29 August 1797), styled Wright of Derby.
A flask in which a large quantity of urine has been boiled down, is seen bursting into light as the phosphorus, which is abundant in urine, spontaneously ignites in air.

By evaporating urine, Brand had produced ammonium sodium hydrogenphosphate, which on further heating produced sodium phosphite. In the presence of carbon (charcoal) this was decomposed to produce white phosphorus and sodium pyrophosphate.

1. (NH4)NaHPO4 → NaPO3 + NH3 + H2O

2. 8NaPO3 + 10C → P4(white) + 2Na4P2O7 + 10CO

In 1680 Robert Boyle and later his assistant Frederick Slare devised improved methods of extracting phosphorus from urine using sand.

3. 4NaPO3 + 2SiO2 + 10C → P4 + 2Na2SiO3 + 10CO

One advantage was that this method liberated all the phosphorus from the sodium phosphite.

The exact nature of the element was not understood for some time. The fact that a minute portion of it, rubbed between two pieces of paper took fire, caused excitement all over Europe whenever it was demonstrated. Given that its isolation was by manipulation of large quantities of urine and that only minute quantities of phosphorus were produced, meant that the substance attracted enormous prices. This only changed in 1769 when the Swedish chemists Carl Scheele and Johan Gahn discovered that phosphorus could be extracted from calcium phosphate (Ca3(PO4)2) found in bones, enabling greater quantities at much cheaper prices to become available.

By the 1850's Albright and Wilson had replaced the raw material for phosphorus production so that instead of animal bones, mineral phosphates were used. The bones or rock were dissolved in sulfuric acid to give phosphoric acid and calcium sulfate as a by-product. The acid was concentrated, mixed with 25% of its mass with carbon, dried in iron pots to a black powder and then distilled in clay retorts. The phosphorus that distilled over was condensed into 25-30 lb. blocks called "cheeses". After refining and casting into sticks (all under water to prevent it catching fire), the product was shipped, again under water, to the end users.

Elemental phosphorus exits in a number of different allotropes.

Military uses of white phosphorus are not covered in this course.

White phosphorus
The most important form of elemental phosphorus from the perspective of applications and the chemical literature is white phosphorus. It consists of tetrahedral P4 molecules, in which each atom is bound to the other three atoms by a single bond. This P4 tetrahedron is also present in liquid and gaseous phosphorus up to the temperature of 800 °C when it starts decomposing to P2 molecules. Solid white phosphorus exists in two forms. At low-temperatures, the β form is stable. At high-temperatures the α form is predominant. These forms differ in terms of the relative orientations of the constituent P4 tetrahedra.

Allotropes of Phosphorus



violet - Hittorf


White phosphorus is the most reactive, the least stable, the most volatile, the least dense, and the most toxic of the allotropes. White phosphorus gradually changes to red phosphorus. This transformation is accelerated by light and heat, and samples of white phosphorus almost always contain some red phosphorus and accordingly appear yellow. For this reason, white phosphorus that is aged or otherwise impure (e.g. weapons-grade not lab-grade white P) is sometimes called yellow phosphorus. White phosphorus glows in the dark (when exposed to oxygen) with a very faint tinge of green and blue, is highly flammable and pyrophoric (self-igniting) upon contact with air and is toxic (causing severe liver damage on ingestion). Owing to its pyrophoricity, white phosphorus has been used as an additive in napalm. The odour of combustion of this form has a characteristic garlic smell, and samples are commonly coated with white "phosphorus pentoxide", which consists of P4O10 tetrahedra with oxygen inserted between the phosphorus atoms and at their vertices. White phosphorus is insoluble in water but soluble in carbon disulfide.

Red phosphorus
In 1847 Anton von Schrotter found that sunlight changed white/yellow into red phosphorus, even when moisture and atmospheric oxygen were rigorously excluded. The red product was separated from the residual yellow phosphorus by treatment with carbon disulfide. Red phosphorus was also prepared from the yellow variety by heating it to about 250 °C. in an inert gas. Heating to higher temperatures reconverted the red modification to the yellow one.
Red phosphorus exists as an amorphous network and does not ignite in air at temperatures below 240 °C.

Violet phosphorus
In 1865, Johann Hittorf heated red phosphorus in a sealed tube at 530 °C. The upper part of the tube was kept at 444 °C. Brilliant opaque monoclinic, or rhombohedral, crystals sublimed.
Hittorf found as well that when phosphorus was recrystallized from molten lead, a red/purple form was obtained. Therefore this form is sometimes known as "Hittorf's phosphorus" (or violet or α-metallic phosphorus).

Black phosphorus
Black phosphorus is the thermodynamically stable form of phosphorus at room temperature and pressure. It is obtained by heating white phosphorus under high pressures (12,000 atmospheres). In appearance, properties and structure it is similar to graphite, being black and flaky, a conductor of electricity, and having puckered sheets of linked atoms.
Black phosphorus has an orthorhombic structure and is the least reactive allotrope: a result of its lattice of interlinked six-membered rings. Each atom is bonded to three other atoms.

History of the Match
The origins of the match have been traced back to Northern China in the 6th Century. The details have been lost but it has been recorded that during a siege on their town the women wanted to conserve fuel and so rather than keeping fires continuously running devised a method of chemically treating sticks such that they could be used to readily light a fire when needed.

It is a long jump from there until the 1770's when white phosphorus was used to coat the tip of a slip of paper and sealed in a glass tube. When a flame was needed, the glass tube was broken; air entered, oxidized the phosphorus and ignited the paper. A modification of this was called the "phosphorus box" which consisted of sulfur-tipped wooden splints sold with a small bottle of phosphorus. To start a fire, the splint was dipped in the phosphorus and then held in air. Air burned the phosphorus, which burned the sulfur and then the wooden splint caught alight.

Friction lights
By 1827 John Walker, a chemist/apothecary from Stockton-on-Tees, was already selling his friction matches in boxes of 100. Walker (29 May 1781-1 May 1859) was a surgeon's apprentice from a rather early age and, during his apprenticeship, conducted many chemical experiments that gave him considerable local status as a scientist. Being of a studious and retiring nature and not caring for the sometimes revolting duties required of a surgeon in those early days, Walker abandoned the profession of surgery and went to work for a firm of druggists.
He dipped wooden splints in sulfur and tipped them with a paste made from potassium chlorate, sugar and antimony trisulfide. The match was ignited by drawing through a fold of sandpaper. Sometimes it didn't work and sometimes a flaming ball flew off and landed on the carpet or a dress. Some reports suggest that Walker discovered the idea of a friction match by accident when some 'percussion powder' (a mixture of potassium chlorate and antimony sulfide) that he was making fell on a stone hearth and ignited. The first recorded sale was in 1827 but he could well have been making them a year earlier. These matches were prohibited in Germany and France because they were thought to be too dangerous.

In 1829 a young chemist named Samuel Jones made and sold imitations of Walker's matches, calling them Lucifers (Walker never patented his idea). Jones, who may have received his inspiration from lectures by Michael Faraday, advertised his goods by stating that they had been "lectured upon" at the London and Royal Institutions. Jones was the first match maker to sell his product in small rectangular cardboard boxes.
Early directions for use, printed on the boxes of Jones' lucifers, included the following caution:
"If possible, avoid inhaling gas that escapes from the combustion of the black composition. Persons whose lungs are delicate should by no means use the Lucifers."
Early Lucifer match composition (1832-1833)
Sulfur 6.5%
Antimony sulfide24.6%
Potassium chlorate27.6%
Ferric oxide5.6%
Gum arabic35.7%

Walker is often credited with inventing the friction match, although neither his nor Jone's matches contained phosphorus. The idea of using friction to generate the heat needed to ignite the match has been followed ever since, rather than using spontaneous chemical reactions to produce a flame.

The first phosphorus match is credited to a Frenchman, Dr. Charles Sauria of St. Lothair. Dr. Sauria made his phosphorus-tipped splints late in 1830, with a formula essentially the same as that used for lucifers, except that phosphorus was used instead of for antimony sulfide.

Impregnation of match heads and splints began around 1870 and attributed to an Englishman named Henry Howse who was granted American patent 123905, 1872 for "safety matches". The flame-retardant chemicals listed in Howse's patent included alum, magnesium sulfate, sodium tungstate and silicate, ammonium borate, chloride, and phosphate, zinc sulfate, and salts of both alkaline and alkaline earth metals.

In 1882 William Pitt obtained US patent 256920 for a match splint that was flame-proofed at only one end. This type of match, commonly referred to as the "drunkard's match," was designed so that the splint would extinguish itself about half-way down, thus protecting the user's fingers from accidental burns. The impregnating material was generally a concentrated solution of sodium silicate or sodium tungstate, coloured with a suitable dye.

Two problems were associated with the original "strike-anywhere" matches. First, white P is extremely toxic (0.1g can be fatal). Phosphorus vapour is oxidised in air to phosphorus(V) oxide (P4O10) and it is the oxide that is dangerous. It is taken into the body through cavities in the teeth and destroys the jaw, causing 'phossy jaw' or phosphorus necrosis. It was recognised that it did not attack people with sound teeth and so the match manufacturers introduced free dental treatment and regular inspections to protect their workers.
Second, the matches were often ignited accidently by rats in warehouses gnawing through the boxes etc.

The obvious replacement was red-phosphorus which is less toxic. However the ignition temperature is much higher (260 °C compared to 30 °C for white phosphorus).

Although the reactions of red and white phosphorus are generally the same, most reactions are carried out using the red allotrope given that it is much safer to handle. Partial combustion of red-P in air produces P4O6,
P4 + 3O2 → P4O6
where the P is in the 3+ oxidation state.

Further oxidation leads to P4O10 (often called phosphorus pentoxide since that was the originally proposed composition) where the P is in the +5 oxidation state.
P4 + 5O2 → P4O10

Nowadays matches are usually pre-treated with a solution of monoammonium phosphate (NH4H2PO4) that prevents afterglow when the flame is extinguished (important given that so many matches have been thrown out of moving vehicles).

The modern "safety-match"
The match head now consists of a paste containing an oxidant such as KClO3 and phosphorus sesquisulfide (P4S3), instead of a mixture or red phosphorus and sulfur. (P4S3 was possibly first suggested in 1872 by Frederick Zaiss of Philadelphia in US Patent 125,874 on coloured and scented parlour matches) These "safety-matches" were consider non-toxic and unaffected by the atmosphere and did not react readily with water. The mixture burnt in air above 100 °C and were held together with glue.

The structure of P4S3
Friction, caused by striking the match on the side of the match box that has powdered glass held on by glue, ignites the match by initiating the exothermic reaction between the P4S3 and the oxidant.

Mass production
On September 11, 1888, Ebenezer B. Beecher of Westville, Connecticut, was issued US Patent 389435 for a continuous, automatic match machine, that revolutionized the match industry and laid the foundation for the modern high-speed continuous machines.

The manufacturing of "safety matches" by Bryant and May, the largest UK match manufacturer began in 1900, although they had been in the match business since 1861. They finally closed down their last plant in Liverpool in 1992.

In the 1950's it was claimed that the USA used around 57 million matches every hour! With the reduction in the number of chain-smokers due to an awareness of the health risks of smoking (and with the advent of cigarette lighters) this number will no doubt have dropped dramatically.

A 1978 US Patent issued to Stanley J. Radkwski, Wilbraham; John M. Lawrence, Monson; Enever Naggar, Longmeadow; Raymond W. Dunham, Springfield, all of Mass. provided details for the production on a self-extinguishing paper match used in the form of book matches.

The general formulation of the flame retardants were:
Weak Flame Retardant: Ammonium Chloride 0-20%
Strong Flame Retardants: Monoammonium Phosphate 0- 5%
Diammonium Phosphate 0- 5%
Remainder: Water

Much of the information in these course notes has been sourced from Wikipedia under the Creative Commons License. Students taking this course will be encouraged to contribute to Wikipedia as a part of their course assignments.

Return to CHEM2402 course outline.

M.F. Crass, A History of the Match Industry, 1941 J. Chem. Educ., 18, part 1 116-120, part 2 277-282, part 3 316-319, part 4 380-384, part 5 428-431
John Emsley, "The 13th Element, The sordid tale of Murder, Fire and Phosphorus", John Wiley and Sons, Inc., New York, 2000.
Yuen Chu Leung, Jurg Waser, S. Van Houten, Aafje Vos, G. A. Wiegers and E. H. Wiebenga, The crystal structure of P4S3, Acta Cryst. (1957). 10, 574.
Peter E. Childs University of Limerick, Limerick, Ireland: Phosphorus: from urine to fire, Pt 1. and Phosphorus: from urine to fire, Pt 2.

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