Phosphorus and the history of the match
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.
(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
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
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.
exits in a number of different allotropes.
Military uses of white phosphorus
are not covered in this course.
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
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.
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 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.
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
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)
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
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.
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.
The structure of P4S3
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%
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
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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