1. Introduction
  2. The Element Phosphorus
  3. The Mineralogy and Production of Phosphorus
  4. Making Fire
  5. Phosphorus on the Farm
  6. Phosphorus at War
  7. References


In 1669 the Hamburg merchant and alchemist Hennig Brandt heated the residue from evaporating urine with powdered charcoal, and condensed the vapor that was evolved into a waxy solid. This solid glowed in the dark, without heat, an astonishing phenomenon. He called the mysterious substance phosphorus, taken directly from the Greek phosphoros, "light-bringer." This was also the name of the planet Venus as morning star, "Lucifer" in Latin. The discovery created quite a stir, and soon nobody was throwing away urine. Travelling alchemists amazed royal audiences, and it was the talk of the time. A normal person excretes about a gram of phosphorus daily.

It gave the name to phosphorescence, which is the nonthermal emission of light after the stimulus has been removed, in distinction to fluorescence. The light Brandt observed was chemiluminescence, consequent to the combination of the phosphorus with atmospheric oxygen to form the trioxide. This was burning, but the light is not due to thermal excitation. When the pentoxide is formed, there is no chemiluminescence.

Phosphorus is an important agricultural fertilizer, an essential element in metabolism and the transfer of biological energy, a component of matches, an ingredient in pyrotechnic applications, and besides of considerable scientific interest. Some of the properties, applications and lore of phosphorus are reviewed in this article. Phosphorus presents many puzzles that are unsolved or only partially solved, which will make this discussion more interesting.

The Element Phosphorus

Phosphorus is in column VA of the periodic table, between nitrogen and arsenic, neither of which it resembles. In the second period, silicon is its neighbor to the left, and sulphur to its right. Phosphorus resembles sulphur more than it does silicon. It never gives up electrons to become a cation, and holds oxygens so tightly that the H comes off of OH easily, and its hydroxides are acidic. Its atomic number is 15, its atomic weight 30.98 (31 is good enough for government work). Its only stable isotope is P31. P32 is a beta-emitter with half-life 14.5 days, often used as a biological tracer.

Its electron configuration is Ne3s23p3, and its first ionization potential is 10.9V. It can share the three p-electrons to make three covalent bonds, exhibiting a valence of +3 or -3 (though it never loses or gains 3 electrons so that it has this formal charge). All five electrons can be used to form covalent bonds as well, with a valence of +5. Phosphorus forms many compounds where the idea of valence does not work well, because of strange structures.

Phosphorus can form a diatomic molecule P2 with a triple bond, like N2, but only at higher temperatures and in the gas phase. It easily forms a tetrahedral P4 molecule that exists in the gas phase and also as a waxy solid and viscous liquid, melting at 44.1°C and boiling at 280°C. The solid has a density of 1.82 g/cc. This form is called white phosphorus (it is actually colorless), and is the form phosphorus usually takes when it is reduced to the pure element. White phosphorus burns readily in air, forming a thick white smoke of phosphorus pentoxide, P2O5. This exists as a dimer in the vapor, which dissociates a little above 1500°C. It is very stable and has an extreme thirst for water, which it will extract even from ethyl alcohol to form ethylene, and from sulphuric acid to form SO3. It is an excellent drying agent, as may be suspected. White phosphorus is extremely toxic. As little as 150 mg is a fatal dose. It is usually kept under water. Those who sniffed its vapor said it had a garlicky odor before they died.

The stable form of phosphorus at room temperature is black or violet phosphorus. Apparently tetrahedral sp3 orbitals are formed, as in carbon or silicon, and the phosphorus forms a giant molecule. Or, it tries to. Doing this reliably to form a uniform crystal like diamond or single crystal silicon seems to be beyond the skills of the phosphorus atom, and only bits and pieces form, connected in random and uneven ways. This material is called red phosphorus. It was discovered in 1845 when white phosphorus was heated gently (240°C) away from air in the presence of iodine as a catalyst. Red phosphorus sublimes at 280°C rather than melting or boiling. If sublimed at low temperatures in a vacuum, the condensate is red phosphorus, but usually the condensate is white phosphorus. Its density is 2.05 to 2.34 g/cc. It is insoluble in anything, and is nonpoisonous. Some sources say it can be noxious under certain conditions, but this does not seem to be borne out by experience, and it can be assumed completely safe so long as you do not make a diet of it. Red phosphorus does not glow in the dark. It is useful for making phosphorus compounds, and in matches.

When red phosphorus burns, the evaporated phosphorus will condense as white phosphorus, which creates a fire (reignition) and health hazard. For this reason, red phosphorus fires should be thoroughly cleaned up. If red phosphorus contains a small amount of neutralizing agent, such as CaCO3, it is almost indefinitely stable. On burning, it produces P2O5 smoke, as white phosphorus does.

White phosphorus usually has a yellow tinge due to small amounts of red phosphorus, so it is sometimes called yellow phosphorus. The red and yellow forms are good examples of allotropes, like diamond and graphite. Red phosphorus was long described as amorphous, but what this term really means is "we don't know its structure" when encountered in the literature.

Like silicon, phosphorus forms bonds with itself only under duress, when there is nothing else around. Also like silicon, it enjoys forming -P-O- bonds, and can even make -O-P-O-P-O-P-O- chains, but not very long ones. This property is essential to life. The DNA and RNA of heredity are built on backbones of sugar molecules (deoxyribose and ribose) connected by -O-P- links. The carrier of energy in the cell, ATP (adenosine triphosphate) contains an O-P-O-P-O-P-O chain that is the key to its utility. ATP is discussed in Burning Food. Phosphorus forms several sulphides, among them P4S3, used as a fuel in pyrotechnic mixtures, such as in the tip of a strike-anywhere match, where it replaces the dangerous white phosphorus.

Phosphorus pentoxide is hydrolyzed in water to form one of several phosphoric acids. With one molecule of water, metaphosphoric acid, HPO3, is formed. With two molecules, pyrophosphoric acid, H4P2O7, is formed. With three molecules, or an excess of water, orthophosphoric acid, H3PO4 appears. In each of these acids, the acidic H+ comes from an OH group attached to the P atom, and the ionization, as usual, is stepwise, with the first hydrogen coming off easily, the last with difficulty. These acids form salts with metallic cations, such as trisodium phosphate, Na3PO4. These salts form alkaline solutions in water, as the phosphate ties up H from the water to free the OH. Trisodium phosphate is called TSP, which is very commonly used as a detergent and water softener. Dishwasher soap is largely TSP.

Phosphorus trioxide, P2O3 is formed by burning phosphorus in a deficiency of air. It is colorless, melting at 23.8°C and boiling at 173°C. It actually exists as the dimer. In water, it hydrolyzes to form orthophosphorous acid, H3PO3. Of the three hydrogens, only two are part of OH groups and ionize. The remaining hydrogen is bonded directly to the P, and stays there. Note the spelling: this is -ous, while the element is -us. The salts of orthophosphorous acid are called phosphites. Orthophosphorous is a stronger acid than orthophosphoric acid. Hypophosphorous acid, H3PO2, has two hydrogens directly bonded to phosphorus, and only one OH group that can ionize. It is an even stronger acid than orthophosphorous. All this is just a glance at the wonderful world of phosphorus oxides and acids.

Phosphorus forms phosphine, PH3, which is more like arsine than ammonia. It is a pyramidal molecule of the expected shape. It does not form by direct union of the elements, but only indirectly. It can be made by boiling white phosphorus in a solution of potassium hydroxide. Oxygen must be purged from the apparatus. If the gas is allowed to bubble up through water, it makes smoke rings of P2O5 on encountering the air, as the phosphine burns. Phosphine does not actually spontaneously burn except in pure oxygen, but some P2H4 produced concurrently does inflame spontaneously, and ignites the phosphine. Phosphine is also produced by the action of water on calcium phosphide, Ca3P2 → 2PH3 + 3Ca++ + 6OH-. Canisters of calcium phosphide are used as signals at sea. When pierced and thrown into the water, they make volumes of smoke, and cannot be extinguished. Phosphine is only slightly soluble in water, and is very poisonous.

Phosphorus forms halides like PCl3 and PCl5. White phosphorus combines directly with chlorine to make these compounds, first the trichloride, and then the pentachloride. The trichloride hydrolyzes to orthophosphorous acid, the pentachloride to orthophosphoric, and hydrogen chloride. The pentachloride will extract oxygen from compounds, making oxychlorides like POCl3, also called phosphoryl chloride.

Although white phosphorus, and many phosphorus compounds, such as phosphine, are toxic, orthophosphates and other compounds are nontoxic, and can be added to food. A "phosphate" was a drink made with fruit syrup, soda water, and phosphoric acid, which gave it an extra tang. Acid and basic phosphate salts are gentler than the carbonate equivalents (because phosphoric acid is stronger than carbonic). Acid phosphate salts (such as primary calcium phosphate, Ca(HPO4)2) are used in baking powder instead of tartaric acid, and basic salts (such as TSP) are used in cleaning powders. Because phosphate is a good fertilizer, the use of phosphates in cleaning agents is now discouraged to avoid eutrophication of the waste water. The "substitute" now on sale is sodium carbonate, which works even better than the phosphate, but dissolves skin. Fudge brownie mix contains tricalcium and monocalcium phosphate, devil's food cake mix sodium acid pyrophosphate and dicalcium phosphate, scalloped potato mix disodium phosphate, and mashed potato flakes sodium acid pyrophosphate, all well down in the lists. Phosphates are not to be feared in the least, and are among the safest additives of all. Monobasic sodium phosphate, NaH2PO4, is called tasteless salt because it looks like salt, but has no salty taste.

The Mineralogy and Production of Phosphorus

The only common mineral of phosphorus is apatite, Ca5F(PO4)3. Apatite is a family of minerals, of which this one, called fluoapatite is the commonest. Its hardness is 5, medium hard, and its density is 3.15-3.20 g/cc. It occurs in two principal forms, crystalline apatite, and phosphorite, which is cryptocrystalline. Crystalline apatite is of inorganic origin, while phosphorite is the remains of animal bones. The largest reserve of crystalline apatite is in the Kola Peninsula, near Murmansk, Russia. The largest reserve of phosphorite is in Idaho and Wyoming in the United States. Wavellite is hydrous aluminium phosphate, a rare mineral usually formed as a secondary deposit from phosphorite. Phosphates are found in a variety of rare minerals, but never in significant quantities. The generic term for a rock from which phosphorus can be economically produced is phosphate rock.

The 8 square mile island of Nauru, to the east of Australia at 167°W and 0.5°S is a plateau of phosphate rock surrounded by beaches. The phosphates are probably the remains of bird droppings, and are the only resource of the island. It is inhabited by about 8,000 Polynesians and a few Chinese, and became independent 31 December 1968. The great phosphate wealth was squandered. After mining ceased in the 1990's, the islanders returned to poverty. There are attempts to restart mining and to find other sources of income.

Elemental phosphorus is produced in electric furnaces by heating phosphate rock, sand and coke. The silica combines with the phosphate rock to give calcium silicate, a slag, and phosphorus pentoxide. The pentoxide is then reduced by the coke to give phosphorus vapor and carbon monoxide. The phosphorus is then condensed and cast into sticks, which are kept under water. The carbon monoxide can be used as a fuel.

Fertilizer, which is a soluble phosphate, is produced by treating phosphate rock with sulphuric acid. More information on phosphate fertilizers is given below, under Phosphorus on the Farm. In the past, most phosphate fertilizer has been made from Florida phosphate rock, which is close to its markets, but this resource is near exhaustion. The largest reserve of phosphorite in the world is in southeastern Idaho and western Wyoming, in the Phosphoria formation of Permian age. The Phosphoria is a widespread formation, consisting of shales, limestones and chert in addition to phosphorite, which is localized in certain areas. It includes oil shales, petroleum, and the black shales that are the source of petroleum. Quite clearly, it was associated with a time of abundant life in shallow seas.

The end of the Permian, when the Phosphoria was laid down, was also the time of the greatest extinction of life that the Earth has known. Nearly 90% of all life perished for unknown reasons in a geologically brief interval, and the Earth made a new start in the Triassic period that followed. The phosphorite, an unusual deposit, probably records a great dying, when the phosphorus from the bones of countless animals was converted to apatite in a shallow basin that was covered by sediment before the deposit could be dispersed.

Bones and teeth are made of apatite, a specially hard version of which forms tooth enamel. If there is insufficient fluorine, then the apatite that is made is faulty, and the teeth are soft and decay easily. Bone is made of hydroxylapatite, but there is fluorapatite in teeth. Many areas of the United States have too little fluorine to make good bones and teeth, so it is added to the drinking water in extremely small amounts that are, nevertheless, sufficient. The usual amount is 1 ppm. Denver water contains an average of 0.91 ppm of fluoride. A level of 4 ppm is the maximum contaminant level of fluoride in drinking water. A good deal more than this is required to cause mottled teeth, seen naturally in some areas with high fluoride. By comparison, the sulphate content was 35.4 ppm, the sodium content 14.0 ppm, and the trihalomethanes (from chlorination) were 40 ppm.

Making Fire

Fire gives us heat for warmth and cooking, light when the Sun is down, smoke for driving away mosquitoes and signalling, eases the hollowing-out of logs to make dugout canoes, and leaves ashes to make alkalis for soap and hominy. Well, we don't do all of those things any more, but lighting fires is still something we do, and when the fires are not automatically lighted for us, we mostly use the convenient wooden match for the purpose. Matches usually use phosphorus, which is the excuse for talking about firemaking here.

The word match comes via French mèche, a "wick," as in a lamp. The match, or more precisely, the slow match, was a piece of rope soaked in potassium nitrate, KNO3, and dried. When lighted and blown out, a red ember continued burning slowly without flame. When a cannoneer wanted to discharge his piece, he touched the match to the touchhole, which had been filled with gunpowder, and the cannon was "fired." A musketeer clipped his match to the matchlock mechanism at the breech of his weapon, and pulled it back against a spring until a catch engaged. Then, when he pulled the trigger, the match was quickly brought to the firing pan, and the firearm discharged. The match, then, was a way of having fire ready for when it was needed, so it is quite reasonable to transfer the name to the wooden match. The quick match was a straw filled with black powder, with a short string that had been coated with sulphur inserted in one end. This was used in blasting.

The lighting of small fires for discharging firearms, lighting pipes and cigarettes, or lighting gas jets, was not required in most of human history. The problem was starting the fire in the hearth, or the campfire, which often burnt continuously for hours. It was not unreasonable to devote some effort and time to the process. First, one needed a spark (we'll discuss how to get sparks below). The spark was caught in tinder, which caused the tinder to make a small glowing ember. Tinder was dry, powdery, easily-combustible matter. Charred linen made a good tinder. Rotten wood could be crumbled into a dry powder. In America, this was called punk. Dry leaves can be powdered to make tinder. Good tinder is essential to fire-lighting. The small ember was then turned onto some fine combustible material, such as dried grass or shavings of resinous pine wood, or finely shredded paper. By gentle blowing, the temperature of the ember could be raised sufficiently to ignite this material, and a visible flame appeared. This fire was then used to light the kindling, usually split sticks, but now often crumpled newspaper, beneath the main combustibles, large pieces of wood or coal that would burn sedately. This effort usually gave great satisfaction, as the cheery fire where before there was only a cold hearth was observed, and the radiated heat warmed the face.

Every traveller carried a tinder-box containing a piece of steel, a piece of flint, and a supply of reliable tinder. When the steel was struck against the flint, a small fragment of steel was dislodged and heated by the distortion of the impact. Finely divided iron catches fire readily in air, so it was a burning spark that landed in the tinder. (This is what I believe to be the mechanism; if this is incorrect, I am eager to know.) These sparks are produced in profusion at a grinding wheel. Gasoline trucks have a dangling chain to discharge static electricity. This chain can be seen striking sparks on a concrete road, which I think is somewhat self-defeating. The little glowing granule of tinder is then tipped out onto the kindling, so the fire can be blown into life. Pyrite also gives sparks when struck by steel, and should give sparks when struck by flint as well, since I believe the sparks in this case are from the pyrite.

Flint and steel are represented at the present day by the steel wheel and "flint" of the cigarette lighter, or a squeeze lighter for an oxyacetylene torch. The "flint" is not flint, but an alloy of 70% mischmetall (a 50:50 alloy of cerium and rare earth metals) and 30% iron, or a spinthoteric lead-zirconium alloy, that gives off copious sparks, as pyrite did in the wheellock, held against the wheel by a spring. The sparks are caught on a wick saturated with a combustible liquid, or pass near an orifice from which combustible gas is issuing. People seem to prefer these to matches, since there is no burnt splint to dispose of, and they burn for as long as required. I made up the word "spinthoteric" from the Greek for "spark-producing," since the common term "pyrophoric" is misused for this property. "Pyrophoric" should be restricted to substances that spontaneously catch fire upon exposure to air.

Indians originally had no steel, so sparks were a rare commodity. A fire was kept going in the village, since fires were so hard to light impromptu. Instead of waiting for lightning, friction could be used to create a glowing fragment. The most efficient way to do this is with a drill operated by a bow. A hard wood rod is rapidly rotated back and forth in a soft wood, which provides the tinder automatically. With luck and effort, an ember is created with which a fire can be started. In popular lore, this is "rubbing two sticks together," which greatly depreciates the skill involved. Most people today could not create a fire from scratch in any way.

It was a generalization, with great substantiation in practice, that the Indian built a small fire adequate to his needs, while the European built a huge roaring blaze one had to stand well back from, and which often escaped to wreak havoc. The Indian built a small fire, and moved without leaving traces, not for any environmental sensitivity, but to avoid the notice of enemies. I suppose the European thought the roaring fire would scare away bears and Indians, but it did not work out that way, merely broadcasting the location so warriors could lurk in the shadows and pick off the European enemies as they wandered out to relieve themselves, or were posted as sentries. The European needed sentries, while the Indian found that the European broacast his approach with noise by day and fire by night. Even the bears might have been attracted by the promise of easy prey.

But back to matches! In 1680, only a decade after the discovery of phosphorus, Robert Boyle made sulphur heads on thin splints, using glue to hold them together. He prepared paper that had been soaked in dissolved phosphorus and then dried. When a piece of the paper was doubled and a sulphur head drawn through it, the sulphur caught fire. These appear to be the first matches, but they were only a curiosity, since the phosphorus paper had to be fresh, and was inconvenient to prepare from the rare substance.

A predecessor to the match was the pyrophorus, "fire-carrier," a sealed tube containing a pyrophoric mixture that would take fire when the tube was broken. Homburg's pyrophorus was a carbonized mixture of alum, flour and sugar, a sort of activated charcoal. Hare's pyrophorus was roasted Prussian blue. Roesling's pyrophorus was a powder to be packed on top of the tobacco in a pipe, and ignited by sucking air through it. Faraday demonstrated a "lead pyrophorus" made from roasted lead tartrate sealed in a glass tube. On breaking the tube, the powder gave a red flash when shaken out.

The "Instant Light Box" appeared around 1780. It contained splints with heads of potassium chlorate, KClO3, a strong oxidant, and a bottle of concentrated sulphuric acid. When a split was dipped into the acid, it instantly caught fire. These matches were something of a commercial success, though the inconveniences were obvious. In 1827, John Walker of Stockton-on-Tees invented the lucifer (the Latin equivalent of "phosphorus"), which was a splint with a head of a mixture of KClO3 and antimony trisulphide, Sb2S3, held together with gum and starch. These ignited (sometimes) on being scratched on sandpaper, without the inconvenient necessity for acid.

A few years later, in 1831, Charles Sauria in France substituted phosphorus sulphide, P4S3, for the antimony sulphide, which made a very reliable match that was easy to ignite. The reaction to prepare the sulphide from white phosphorus is vigorous, and phosphorus vapor poisoned the workers. Their disease was the terrible "phossy jaw" with necrosis of the jawbone, the principal symptom of chronic phosphorus poisoning. It is probably not true that white phosphorus was ever used directly in matches that were sucked by the match workers. However, lead dioxide, PbO2 was used as an oxidizing agent. A white phosphorus--lead dioxide match would have been a remarkable poison.

In 1845, A. Schrötter of Vienna discovered how to prepare red phosphorus, which is insoluble and nonpoisonous. Phosphorus sulphide could now be prepared safely from red phosphorus. Today's matches are made with pine sticks impregnated with ammonium phosphate to discourage afterglow, and impregnated with paraffin wax for easy burning. The heads of "strike-anywhere" matches are a mixture of potassium chlorate and sulphur or rosin, held together with glue. On the tip is a mixture of P4S3 and pulverized glass for friction.

The "safety match" is the most familiar pyrotechnic device. It is reliable, and, as its name indicates, safe. It consists of two parts, the paraffin-impregnated splint with a head that constitutes the "first fire" of the device, and a striking strip. The head contains 45-55% KClO3, with a little sulphur and starch, a neutralizer (ZnO or CaCO3), 20-40% of siliceous filler, diatomite and glue. Some heads contain antimony trisulphide so they burn more vigorously. The striking strip is 50% red phosphorus, 5% neutralizer, 4% carbon black, 25% powdered glass for roughness, and 16% binder. Antimony trisulphide is sometimes also in the striking strip. Friction unites the chlorate and the phosphorus, initiating the reaction, which spreads to the fuels in the head, and from there to the splint. None of the components are poisonous or dangerous. A match head could explode if struck with a hammer on a hard surface, however, but this would not be very dangerous. Moisture is the enemy of matches, but the components will recover their activity on drying.

Potassium chlorate and red phosphorus are very safe when used as in safety matches, but mixing the pure compounds is extremely hazardous.

Phosphorus on the Farm

Phosphorus is essential to life. No animal or plant can live without it. In addition to using it in metabolism, many animals use it in their skeletons. About 60% of bone is calcium phosphate, 40% calcium carbonate. Animals must have phosphate in their diets, while plants must find it in the minerals they obtain from the soil. The central role of phosphorus is sufficient excuse for mentioning it here.

The 14 elements necessary for plant life are: C, O, H, N, Ca, P, K, Mg, Fe, S, B, Mn, Cu and Zn. C, O and H are supplied by the air and water. The remainder must come from the soil. The elements that must most usually be supplemented in farming soil are N, P and K. Nitrogen and potassium are easily leached out by water, and phosphorus is removed in the crops. Grasses, cereal grains and beans take large quantities of phosphorus. These are used to feed animals that are sold off the farm, carrying phosphorus with them, or are sold directly. Natural replenishment of phosphorus is very slow, since it is a scarce mineral, so all farming land, arable or grazing, becomes depleted in time. No amount of "organic" farming can make this up. Farm manure, in particular, is very poor in phosphorus. Land depleted in phosphorus will not produce food.

The ancient way to overcome phosphorus deficiency was to spread crushed animal bones over the land. Now, phosphate rock is mined and treated with sulphuric acid to make superphosphate, through the reaction Ca3(PO4)2 + 2H2SO4 + 5H2O → Ca(H2PO4)3·2H2O + 2(CaSO4·2H2O). The soluble acid phosphate and the gypsum are spread on the land together. By using phosphoric acid made from one portion of the rock to treat another portion, a molecular ratio of 1:1 can be obtained rather than 1:2, so there is more phosphorus per pound. This is called treble superphosphate. A fertilizer is rated according to the soluble phosphate, or available phosphoric acid. Superphosphate contains 14% to 20% available phosphate, the treble superphosphate 40% to 50%. Virtually everything with some phosphorus in it has been used for fertilizer.

Atmospheric nitrogen is useless for plants. It must be "fixed," or converted to soluble form, before it can be used. This can be done by bacteria that live symbiotically in root nodules on legumes. Green manures are legumes specially grown for their nitrogen, such as clover. Alfalfa is specially efficient in adding nitrogen, and makes valuable forage at the same time. Crop rotation with alfalfa can triple the yields of cereal crops. Green manures also supply organic matter, which is also important for soil health. Rye and buckwheat are non-legumes grown for this purpose. Rye is winter-hardy and can be planted after other crops are harvested. Each ton of dry matter in clover turned over into the soil provides 120 lb of nitrogen, from the tops and roots.

Nitrogen stimulates vegetative growth, so too much can as easily be used as too little. Nitrogen deficiency produces stunted growth and a yellowish color, since nitrogen stimulates the production of chlorophyll. Artificial nitrogen fertilizer can be applied as nitrates or ammonia. Nitrates are rapidly leached away, so ammonia is preferable. Potassium is easily leached out of sandy, light soils, and is locked up in insoluble form in heavy soils, so supplementation is often required. Potassium is applied as KCl, K2SO4, KNO3, seaweeed or wood ashes. Alfalfa uses great amounts of potassium, as do sugar beets, corn, potatoes and tobacco. Commercial fertilizers usually supplement nitrogen, phosphorus and potassium. A fertilizer containing all three is called complete. They are characterized by three numbers, which are the percentages of nitrogen, available phosphoric acid, and soluble potash (KOH). A 4-16-4 fertilizer contains 4% nitrogen, 16% available phosphate, and 4% soluble potash. My source does not give the forms that are the basis for this calculation.

Maintenance of the productivity of farms depends on soil conservation, the application of inorganic fertilizers, and the development of productive and robust strains of plants and animals. The present population of the earth has far outstripped the natural productivity of the land, and depends on agricultural technology to support its existence. The collapse of the sustaining power of the soil seems already to be taking place in Africa, which depends more and more on imports of food from the rest of the world, though the population is not dense by the usual standards.

Phosphorus at War

White phosphorus is both a smoke producer and a particularly nasty incendiary agent, known as WP. Its white smoke has the highest total obscuring power (TOP) of any smoke. It was widely used in World War I in grenades and trench mortar rounds to screen troop movements. Most military smokes are now of other types, often colored with dyes. The 4.2-in. "Chemical" mortar of World War II was developed to throw white phosphorus shells, as well as whatever other chemical or biological agents might be required, but was later also found valuable as a general heavy mortar. This was a simple, light, portable weapon of great power, equivalent to a 105 mm howitzer, but of lesser range. It consisted of a tube about 5 ft long, a steel baseplate, and a bipod support with screws for elevation and traverse. The cylindrical round was simply dropped down the tube, and it sailed away on a high trajectory. The phosphorus sticks to whatever it hits, burns, and if what it has hit is combustible, sets it on fire. White phosphorus burns quickly and cooly and so is not a very effective incendiary agent. It is generally mixed with rubber or polystyrene to slow down the burning. Water will put out white phosphorus temporarily, but as soon as the phosphorus has access to air, it will start burning again. White phosphorus wounds are very unpleasant, since the phosphorus must be thoroughly washed out with a nonpolar solvent that is also noninflammable, for obvious reasons, before the burn can be treated. Carbon tetrachloride would be suitable, but it is dangerous because of the cancer hazard.


W. N. Jones, Inorganic Chemistry (Philadelphia: Blakiston, 1949). Chapter 20. Better descriptive material than in more modern texts.

H. C. Rather, Field Crops (New York: McGraw-Hill, 1942). Chapter IV. A book that I acquired by accident that has proved very informative.

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Composed by J. B. Calvert
Created 17 December 2002
Last revised 15 March 2008