


The complicated characteristics of the chemical industry







Economic aspects






Heavy inorganic chemicals

Sodium carbonate and other alkalies

In 1775 the French Academy of Sciences offered an award for a practical method for converting common salt, sodium chloride, into sodium carbonate, a chemical needed in substantial amounts for the manufacture of both soap and glass. Nicolas Leblanc, a surgeon with a bent for practical chemistry, invented such a process. His patron, the duc d’Orléans, set up a factory for the process in 1791, but work was interrupted by the French Revolution. The process was not finally put into industrial operation until 1823 in England, after which it continued to be used to prepare sodium carbonate for almost 100 years.

The Leblanc process

The first step in the Leblanc process was to treat sodium chloride with sulfuric acid. This treatment produced sodium sulfate and hydrogen chloride. The sodium sulfate was then heated with limestone and coal to produce black ash, which contained the desired sodium carbonate, mixed with calcium sulfide and some unreacted coal. Solution of the sodium carbonate in water removed it from the black ash, and the solution was then crystallized. From this operation derives the expression soda ash that is still used for sodium carbonate.

It was soon found that when hydrogen chloride was allowed to escape into the atmosphere, it caused severe damage to vegetation over a wide area. To eliminate the pollution problem, methods to convert the dissolved hydrogen chloride to elemental chlorine were developed. The chlorine, absorbed in lime, was used to make bleaching powder, for which there was a growing demand.

Because calcium sulfide contained in the black ash had a highly unpleasant odour, methods were developed to remove it by recovering the sulfur, thereby providing at least part of the raw material for the sulfuric acid required in the first part of the process. Thus the Leblanc process demonstrated, at the very beginning, the typical ability of the chemical industry to develop new processes and new products, and often in so doing to turn a liability into an asset.

The ammonia-soda (Solvay) process

The Leblanc process was eventually replaced by the ammonia-soda process (called the Solvay process), which was first practiced successfully in Belgium in the 1860s. In this process, sodium chloride as a strong brine is treated with ammonia and carbon dioxide to give sodium bicarbonate and ammonium chloride. The desired sodium carbonate is easily obtained from the bicarbonate by heating. Then, when the ammonium chloride is treated with lime, it gives calcium chloride and ammonia. Thus, the chlorine that was in the original sodium chloride appears as calcium chloride, which is largely discarded (among the few uses for this compound is to melt snow and ice from roads and sidewalks). The ammonia thus regenerated is fed back into the first part of the process. Efficient recovery of nearly all the ammonia is essential to the economic operation of the process, the loss of ammonia in a well-run operation being no more than 0.1 percent of the weight of the product.

Electrolytic process

Later in the 19th century the development of electrical power generation made possible the electrochemical industry. This not clearly identifiable branch of the chemical industry includes a number of applications in which electrolysis, the breaking down of a compound in solution into its elements by means of an electric current, is used to bring about a chemical change. Electrolysis of sodium chloride can lead to chlorine and either sodium hydroxide (if the NaCl was in solution) or metallic sodium (if the NaCl was fused). Sodium hydroxide, an alkali like sodium carbonate, in some cases competes with it for the same applications, and in any case the two are interconvertible by rather simple processes. Sodium chloride can be made into an alkali by either of the two processes, the difference between them being that the ammonia-soda process gives the chlorine in the form of calcium chloride, a compound of small economic value, while the electrolytic processes produce elemental chlorine, which has nearly innumerable uses in the chemical industry, including the manufacture of plastic polyvinyl chloride, the plastic material produced in the largest volume. For this reason the ammonia-soda process, having displaced the Leblanc process, has found itself being displaced, the older ammonia-soda plants continuing to operate very efficiently but no new ammonia-soda plants being built.

Other important processes

The need for sodium carbonate in the manufacture of soap and glass that led to the Leblanc process also led to the creation of the alkali industry and the chlor-alkali industry, another of the historic landmarks of the chemical industry (see Chlorine).

Sulfuric acid

Sulfuric acid is by far the largest single product of the chemical industry. The chamber process for its preparation on the scale required by the Leblanc process might be regarded as the most important long-term contribution of the latter.

Chamber process

When sulfur is burned in air, sulfur dioxide is formed, and this, when combined with water, gives sulfurous acid. To form sulfuric acid, the dioxide is combined with oxygen to form the trioxide, which is then combined with water. A technique to form the trioxide, called the chamber process, developed in the early days of the operation of the Leblanc process. In this technique the reaction between sulfur dioxide and oxygen takes place in the presence of water and of oxides of nitrogen. Because the reaction is rather slow, sufficient residence time must be provided for the mixed gases to react. This gaseous mixture is highly corrosive, and the reaction must be carried out in containers made of lead.

Contact process

Lead is a material awkward to use in construction, and the process cannot deliver acid more concentrated than about 78 percent without special treatment. Therefore, the chamber process has been largely replaced by the contact process, in which the reaction takes place in a hot reactor, over a platinum or vanadium compound catalyst, a substance that increases the speed of the reaction without becoming chemically involved.


Of the large world production of sulfuric acid, almost half goes to the manufacture of superphosphate and related fertilizers. Other uses of the acid are so multifarious as almost to defy enumeration, notable ones being the manufacture of high-octane gasoline, of titanium dioxide (a white pigment, also a filler for some plastics, and for paper), explosives, rayon, the processing of uranium, and the pickling of steel.

Sources of sulfur

Because sulfuric acid is indispensable to so many industries, its primary raw material is of the greatest importance. The needed sulfur is obtainable from a number of sources. Originally, sulfur came chiefly from certain volcanic deposits in Sicily. By the beginning of the 20th century this source was insufficient, but the supply was augmented by sulfur that occurs underground in the southern United States. This sulfur is not mined but is recovered by the so-called Frasch process, in which the sulfur is melted underground by hot water and the mixture brought to the surface in liquid form.

Other sources of sulfur include the ore iron pyrite, an iron-sulfur compound that can be burned to produce sulfur dioxide, and some natural gases, called sour gas, that contain appreciable quantities of hydrogen sulfide. Certain metal sulfides, such as those of zinc and copper, are contained in the ores of those metals. When these ores are roasted, sulfur dioxide is given off. Sulfur is usually shipped in its elemental form rather than in the form of sulfuric acid.

Under some circumstances, the sulfuric acid stage of manufacture can be avoided. Ammonium sulfate, a fertilizer, is normally made by causing ammonia to react with sulfuric acid. In many parts of the world, abundant supplies of calcium sulfate in any of several mineral forms can be used to make the ammonium sulfate by combining it with ammonia and water. This process brings the sulfur in the calcium sulfate deposits into use. Because deposits of calcium sulfate throughout the world are extensive, development of such a process would make the available resources of sulfur almost limitless.

The sulfur present in low percentages in fossil fuels is a notorious source of air pollution in most industrial countries. Removal of sulfur from crude oil adds to the sulfur supply and reduces pollution. It is less easy to remove the sulfur directly from coal.

Carbon disulfide



Fertilizers represent one of the largest market commodities for the chemical industry. A very large industry in all industrialized countries, it is a very important one for introduction as early as possible into developing countries.

The crucial elements that have to be added to the soil in considerable quantities in the form of fertilizer are nitrogen, phosphorus, and potassium, in each case in the form of a suitable compound. These are the major fertilizer elements, or macronutrients. Calcium, magnesium, and sulfur are regarded as secondary nutrients; and it is sometimes necessary to add them. Numerous other elements are required only in trace quantities; certain soils may be deficient in boron, copper, zinc, or molybdenum, making it necessary to add very small quantities of these. As a great industry, however, fertilizers are based on the three elements mentioned above.

Nitrogen is present in vast quantities in the air, making up about 78 percent of the atmosphere. It enters the chemical industry as ammonia, produced through fixation of atmospheric nitrogen, described below. For phosphorus and potassium, it is necessary to find mineral sources and to convert them into a form suitable for use. These three elements are not used in fertilizer only, however; they have other uses and interact with other facets of the chemical industry, making a highly complicated picture. A schematized overview of some of these interactions is presented in Figure 1.


The simplest part of this diagram is the portion representing potassium. The element potassium is seventh in order of abundance in the Earth’s crust, about the same order as sodium, which it resembles very closely in its properties. Although sodium is readily available in the sodium chloride in the ocean, most of the potassium is contained in small proportions in a large number of mineral formations, from which it cannot be economically extracted. When the use of potassium salts as fertilizers began in the second half of the 19th century, it was believed that Germany had a monopoly with the deposits at Stassfurt, but many other workable deposits of potassium salts were later found in other parts of the world. World reserves are adequate for thousands of years, with large deposits in the Soviet Union, Canada (Saskatchewan), and Germany (East and West).

Potassium chloride is the principal commercial form of potash, and some potassium nitrate is also produced. About 90 percent of the production of these goes to fertilizers. For other purposes, the similar sodium salts are cheaper, but for a few special uses potassium has the advantage. Some ceramic uses require potassium, and potassium bicarbonate is more effective than sodium bicarbonate in extinguishing fires.


Phosphorus presents a more complicated picture. It has many uses other than in fertilizers. By far the largest source is phosphate rock, although some use is made of phosphatic iron ore, from which the phosphorus is obtained as a by-product from the slag. As with potassium, there are extensive reserves. The largest deposits are in North Africa (Morocco, Algeria, Tunisia), the United States (largely Florida), and the Soviet Union, but there are also sizable deposits in numerous other countries.

Phosphate rock is found in deposits of sedimentary origin, laid down originally in beds on the ocean floor. The rock consists largely of the insoluble tricalcium phosphate, together with some other materials, including some fluorine. To be used as a fertilizer, phosphate must be converted to a form that is soluble in water, even if only slightly so.

Phosphoric acid (H3PO4) has three hydrogen atoms, all of which are replaceable by a metal. Tricalcium phosphate, in which all three of the hydrogen atoms are replaced by calcium, must be converted to the soluble form, monocalcium phosphate, in which only one hydrogen atom is replaced by calcium. The conversion is done by sulfuric acid, which converts the phosphate rock to superphosphate, widely used as fertilizer. This operation requires large tonnages of sulfuric acid.

The fertilizer industry is not only a matter of manufacturing the right chemical but also of distribution, getting the right material to the right place at the right time. Fertilizers are made centrally but must be distributed over a large agricultural area. A fertilizer factory is, typically, a large installation, characterized by enormous storage silos; the product is manufactured all the year round, but it requires considerable space to store it until the few weeks during which it is distributed on farmlands.

The weight of the superphosphate is greater than that of the original phosphate rock by the amount of the sulfuric acid added; the superphosphate also carries the dead weight of the calcium sulfate that is formed in the manufacturing process. This dead weight can be reduced by replacing sulfuric acid with phosphoric acid (itself obtained by the action of sulfuric acid on phosphate rock, followed by separating the products; or else by an electric furnace process). This process results in triple superphosphate, in which all the calcium originally in the phosphate rock appears as calcium monophosphate. The useful content of the fertilizer, expressed as the percent of phosphoric oxide, is increased from 20 percent in ordinary superphosphate to about 45 percent in the triple variety, resulting in a better than twofold reduction in the amount of material that must be distributed to provide a given amount of the useful oxide.

Instead of using either sulfuric or phosphoric acid to treat the phosphate rock, nitric acid can be employed. One of the resulting products, calcium nitrate, is itself a fertilizer, so what is obtained is one of the many varieties of mixed fertilizers. Instead of neutralizing phosphoric acid with calcium, which contributes nothing but dead weight, ammonia can be used, giving ammonium phosphate, in which both constituents contribute fertilizer elements. Such improvements in fertilizers are constantly being made.

Many other compounds of phosphorus are used. One group is composed of phosphoric acid and various phosphates derived from it. The acid itself is used in soft drinks for its pleasant taste when sweetened and its nutritive value. Other food applications include the use of disodium phosphate in processed cheese; and phosphates in baking powder, flameproofing, and the treatment of boiler water in steam plants. An important use of some of the phosphates is in detergents, discussed below.

Elemental phosphorus exists in many allotropic forms. White phosphorus is used in rodent poison and by the military for smoke generation. Red phosphorus, comparatively harmless, is used in matches. Ferrophosphorus, a combination of phosphorus with iron, is used as an ingredient in high-strength low-alloy steel. In addition, the many organic compounds of phosphorus have varied uses, including those as additives for gasoline and lubricating oil, as plasticizers for plastics that otherwise would be inconveniently rigid, and, in some cases, as powerful insecticides, related to nerve poisons.


The production of nitrogen not only is a major branch of the fertilizer industry, but it opens up a most important segment of the chemical industry as a whole.

Farm manure long supplied enough nitrogenous fertilizer for agriculture, but late in the 19th century it was realized that agriculture was outgrowing this source. A certain amount of ammonium sulfate was available as a by-product of the carbonization of coal, and the large deposits of sodium nitrate discovered in Chile helped for a time. The long-range problem of supply, however, was not solved until just before World War I, when the research of Fritz Haber in Germany brought into commercial operation the method of ammonia synthesis that is used, in principle, today. The immediate motivation for this great development was Germany’s need for an indigenous source of nitrogen for military explosives. The close interrelation between the use of nitrogen for fertilizers and for explosives persists to this day.

Because air is 78 percent nitrogen, there is a little more than 11 pounds of nitrogen over every square inch of the earth’s surface. Nitrogen, however, is a rather inert element; it is difficult to get it to combine with any other element. Haber succeeded in getting nitrogen to combine with hydrogen by the use of high pressure, moderately high temperatures, and a catalyst.

The hydrogen for ammonia (NH3) is usually obtained by decomposing water (H2O). This process requires energy, in some cases supplied by electricity, but more often from fossil fuels. In some cases the hydrogen is obtained directly from the fossil fuel, without decomposing water.

Haber used coke as a fuel. Carbon can burn either to carbon dioxide or, if the supply of air is kept short, to carbon monoxide, by a process known as the producer gas reaction. The gaseous product is a mixture of carbon monoxide with the nitrogen that was originally in the air.

The red-hot coke can also be heated with steam to yield carbon monoxide and hydrogen, a mixture known as water gas. It is also possible to carry out a water-gas shift reaction by passing the water gas with more steam over a catalyst, yielding more hydrogen, and carbon dioxide. The carbon dioxide is removed by dissolving it in water at a pressure of about ten atmospheres; it can also be utilized directly, as noted below. Starting then from water gas, and converting a certain proportion of the carbon monoxide to carbon dioxide and hydrogen, it is possible to arrive at a mixture of carbon monoxide and hydrogen in any proportion.

Synthesis gas

In Figure 1, the words synthesis gas have been shown as the source of two products, ammonia and methanol. It is not quite the same synthesis gas in the two cases, but they are closely related. The mixture of carbon monoxide and hydrogen described above is the synthesis gas that is the source of methanol. But ammonia requires nitrogen, which is obtained from the producer gas by causing it to undergo the water-gas shift reaction, yielding hydrogen. Ammonia requires much more hydrogen, which is obtained from water gas subjected to the water-gas shift. And so, by appropriate mixing, ammonia synthesis gas of exactly the right composition can be obtained.

The above description is a simplified account of how synthesis gas, either for ammonia or for methanol, is obtained from fossil fuel as a source of energy, but it gives an idea of the versatility of the operations. There are many possible variations in detail, depending largely on the particular fuel that is used. The nitrogen industry, which has grown steadily since shortly after World War I, was originally based largely on coke, either from coal or lignite (brown coal). There has been a gradual change to petroleum products as the fossil fuel. As is true with many other branches of the chemical industry, the latest trend is to move to natural gas.

The carbon dioxide removed during the preparation of the synthesis gas can be caused to react with ammonia, often at the same plant, to form urea, CO(NH2)2. This is an excellent fertilizer, highly concentrated in nitrogen (46.6 percent), and also useful as an additive in animal feed to provide the nitrogen for formation of meat protein. Urea is also used for an important series of resins and plastics by reaction with formaldehyde, derived from methanol.

Ammonia can be applied as a fertilizer in numerous forms, ranging from the application of liquid ammonia beneath the surface of the soil, or solutions of ammonia in water (also containing other fertilizer ingredients), or as ammonium nitrate, or other products from nitric acid, which itself is derived from ammonia. Ammonia also has other uses within the chemical industry. The small amount of ammonia consumed in the course of making sodium carbonate by the ammonia-soda process formerly amounted to a considerable volume. Ammonia is used in one process for making rayon, as a refrigerant in large commercial refrigeration establishments, and as a convenient portable source of hydrogen. Hydrogen can be compressed into cylinders, but ammonia, which forms a liquid on compression, packs far more hydrogen into the same volume; it is decomposed by heat into hydrogen and nitrogen; the nitrogen is used to provide an inert atmosphere for many metallurgical operations.

Nitric acid

By far the most important use of ammonia within the chemical industry is to produce nitric acid (HNO3). Nitrogen and oxygen can be made to combine directly with one another only with considerable difficulty. A process based on such a direct combination, but employing large quantities of electrical power, was in use in the 1920s and 1930s in Norway, where hydroelectric power is readily available. It has not proved economical in modern conditions.

Ammonia burns in air, or in oxygen, causing the hydrogen atoms to burn off, forming water and leaving free nitrogen. With the aid of a catalyst, platinum with a small percentage of the related metal rhodium, ammonia is oxidized to oxides of nitrogen that can be made to react with water to form nitric acid.

Nitric acid treated with ammonia gives ammonium nitrate, a most important fertilizer. Ammonium nitrate, moreover, is also an important constituent of many explosives. Three fundamental explosive materials are obtained by nitrating (treating with nitric acid, often in a mixture with sulfuric acid): cellulose, obtained from wood, gives cellulose nitrate (formerly called nitrocellulose); glycerol gives glyceryl trinitrate (formerly called nitroglycerin); and toluene gives trinitrotoluene, or TNT. Another explosive ingredient is ammonium picrate, derived from picric acid, the relationship of which appears more clearly in its systematic name, 2,4,6-trinitrophenol.

A minor but still important segment of the explosives industry is the production of detonating agents, or such priming compositions as lead azide [Pb(N3)2], silver azide (AgN3), and mercury fulminate [Hg(ONC)2]. These are not nitrates or nitro compounds, although some other detonators are, but they all contain nitrogen, and nitric acid is involved in their manufacture.

Related to the explosives are the rocket propellants. A rocket-propelled missile or spacecraft launch vehicle must carry both reactive components (fuel and oxidizer) with it, either in different molecules or in the same molecule. Essentially, rocket propellants consist of an oxidant and a reductant. The oxidant is not necessarily a derivative of nitric acid but may also be liquid oxygen, ozone (O3), liquid fluorine, or chlorine trifluoride.

Other uses for nitric acid not related to explosives or propellants include the production of cellulose nitrate for use in coatings. Without a pigment it forms a clear varnish much used in furniture finishing. Pigmented, it forms brilliant shiny coatings referred to as lacquers. At one time a fibre similar to rayon was made from cellulose nitrate.

Nitrating benzene (Figure 2) yields nitrobenzene, which can be reduced to aminobenzene, better known as aniline. Aniline can also be made by reacting ammonia with chlorobenzene, obtained from benzene. Benzene and ammonia are required in either case. Similar treatment applied to naphthalene (C10H8) results in naphthylamine. Both aniline and naphthylamine are the parents of a large number of dyes, but today synthetic dyes are usually petrochemical in origin (see the article dye). Aniline, naphthylamine, and the other dye intermediates lead also to pharmaceuticals, photographic chemicals, and chemicals used in rubber processing.

The above account gives an idea of the importance, not only for fertilizers but for many other products, of the process known as fixation of atmospheric nitrogen—that is, taking nitrogen from the air and converting it into some form in which it is usable. The tremendous increase in the production of fertilizers has led to the erection of huge ammonia plants. The plants must have available a source of fossil fuel, but petroleum and natural gas are easily transported, so that there is a tendency to locate the plants near the ultimate destination of the product.

A typical plant comprises all the equipment for the preparation of the synthesis gas on the requisite scale, along with equipment for purifying the gas. In the synthesis of ammonia (but not of methanol) any compound of oxygen is a poison for (reduces the effectiveness of) the catalyst, and so traces of carbon dioxide and carbon monoxide must be carefully removed. Compressing the gas to the desired pressure requires extensive engineering equipment. The higher the pressure the greater the yield but the higher the actual cost of compression. The higher the temperature the lower the yield, but the temperature cannot be lowered indefinitely to obtain better yields because lower temperatures slow down the reaction. The temperature used is of the order of 500° C (930° F). The choice of temperature and of pressure is a carefully worked out compromise to give optimal results. The yield equals the amount of nitrogen and hydrogen that combine to form ammonia in any one pass through the converters. Only a fraction is converted each time, but after each pass the ammonia is removed and the remaining gas is recycled. Atmospheric nitrogen contains about 1 percent argon, a totally inert gas, which must be removed from time to time so that it does not build up in the system indefinitely. There is also usually a nearby nitric acid factory and equipment for producing ammonium nitrate in the exact grain size for convenient application as fertilizer.

The growing plants of agricultural crops do not receive all of their nitrogen from synthetic fertilizer. A certain proportion is supplied by natural means. Some plants, notably beans, have a symbiotic relationship with nitrifying bacteria that are able to “fix” the nitrogen in the air and to combine it into a form available for plant life. This natural synthesis takes place without the necessity for pressures of several hundred atmospheres or of high temperatures. In many laboratory syntheses of natural products, great success has been obtained by quiet reactions, without extreme conditions, by a process of following nature gently, rather than by brute force. Ammonia may some day be synthesized by some process that more nearly resembles the natural way; research is being carried out along these lines, and it may be that a far easier approach to fixation of atmospheric nitrogen will be part of the chemical industry of the future.

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