Industrial Inorganic Chemistry: The Alkali Industry

James Richard Fromm

The beginning of the nineteenth century marks the start of industrial chemistry as it is now known. The nineteenth century was the Industrial Revolution of inorganic chemistry. The Industrial Revolution for the chemistry of organic compounds, which occurred near the end of the nineteenth century and accelerated through our own twentieth century, is covered elsewhere. The following Table is intended to give a historical and physical perspective on current chemical production in a large integrated industrial economy. The chemical economies of Canada, and of individual European countries, are more significantly affected by international trade in chemical commodities.

Table: Inorganic Industrial Chemical Production in the United States
		Compound       Production      Production     Major Uses
		               (Tg,1985)       (Tg,1975)
		H2SO4          35.97           29.94          acids
		N2(g)          21.53           19.63          NH3
		O2(g)          15.76           14.61          steelmaking
		NH3            15.71           14.94          fertilizers
		CaO            14.33           17.28          bases, cement
		NaOH            9.88            9.20          paper
		H3PO4           9.54            6.85          fertilizers
		Cl2(g)          9.43            9.13          paper, plastics
		Na2CO3          7.80            6.76          glass
		HNO3            6.68            6.84          NH4NO3
		NH4NO3          6.15            6.52          fertilizer
		CO(NH2)2        6.06            3.50          fertilizer
		CO2(g)          4.20            1.83          CO(NH2)2
		HCl             2.54            1.86          varied
		(NH4)2SO4       1.90            1.59          fertilizer
		C (black)       1.17            1.37          tires

Table Notes: one Tg is also one million metric tonnes (MMT). Minerals which occur naturally and are used but not processed by the chemical industry, such as NaCl, KCl, CaSO4, and S, are excluded from this list.

The literal foundation stone of the alkali industry is chalk or limestone, CaCO3. Beds of limestone are abundant and are found throughout the world. They have arisen over geologic ages from the shells of small marine animals deposited on the beds of ancient seas. Upon strong heating, limestone loses carbon dioxide to give unslaked lime, usually called simply lime, which is CaO:

CaCO3(s) rarrow.gif (63 bytes) CaO(s) + CO2(g)

The construction material cement is made by roasting together a powdered mixture of limestone, sand (silica, SiO2), a clay or shale containing aluminum, and iron oxide at a temperature of up to 1150 K. Modern processing plants use a rotating kiln. The materials lose water and carbon dioxide and partially fuse together; the resulting mass or "clinker" is ground to a fine powder. Often a small amount of gypsum, CaSO4, is added.

The composition of a modern Portland cement is 60 to 67% CaO, 17 to 25% SiO2, 3 to 8%

Al2O3, and up to 6% Fe2O3; small amounts of MgO, MgSO4, Na2O, and K2O may also be present. This cement, when added to sand and crushed stone and mixed with water, gives the artificial stone concrete. The reactions which occur as concrete sets in the presence of water and of atmospheric carbon dioxide are complex. They continue for many years after the initial solidification of the concrete, which slowly continues to harden.

At the end of the eighteenth century, alkali was a general name meaning both potash (potassium carbonate, K2CO3) obtained in crude form from wood ashes, and soda, sometimes called soda ash (sodium carbonate, Na2CO3). Alkali was then in demand for three uses: manufacture of glass, manufacture of soap, and dyeing of fabrics. The primary source of alkali at this time was wood ashes, which on leaching (extraction of soluble material by water) gave solutions of K2CO3. The solid could be obtained by evaporating the water. Soda was obtained from deposits of native soda, either natron (Na2CO3) or trona (sodium sesquicarbonate, Na2CO3.NaHCO3.2H2O), in the same manner.

For glassmaking, the best alkali available was required to obtain a glass of good optical quality. A suitable alkali could be made by strongly heating (calcining) soda, which converts it to the oxide Na2O, after purification of the leach liquid by repetitive filtration and evaporation. The alkali required for soap could be more easily prepared by treating moistened potash with unslaked lime (CaO) to give lye (KOH). Alternatively, treatment of moistened soda with lime gave caustic, or caustic soda (NaOH). These strong bases were used to form soap by boiling fat, olive oil, or tallow together with them. The quality of soap produced depended upon the quality of the ingredients, and was highly variable. Glass, made from potash or soda and available sand, also varied widely in quality depending upon the quality of the ingredients and the skill of the maker. In the absence of commercial sources of pure materials, preparation of glass of optical quality was an art requiring great skill.

Production of Sodium Carbonate

The first major inorganic chemical process developed was the Leblanc process (1787) for the production of artificial soda. The Leblanc process reached significant commercial use in the 1820's.

Example. The first step of the Leblanc process was reaction of salt with sulfuric acid:

2NaCl(s) + H2SO4(l) rarrow.gif (63 bytes) Na2SO4(s) + 2HCl(g)

The resulting sodium sulfate was mixed with chalk and charcoal and heated in a crucible. The 'black ash' product containing Na2CO3 was extracted with water and the pure soda recovered from the solution by evaporation:

Na2SO4 + CaCO3 + C rarrow.gif (63 bytes) Na2CO3 + CaSO4

The waste products, HCl gas and CaSO4 solid, were problems with the Leblanc process. That of the disposition of CaSO4 was never satisfactorily solved, and mountains of leftover calcium sulfate still exist in Europe and Britain. However, partially as a result of the first anti-chemical pollution legislation (the Alkali Act of Great Britain), the HCl gas was soon absorbed by water in towers and sold as a valuable by-product. Heating of a solution of HCl in water first boils off the water, until a concentration of 20.2% HCl by weight is reached. Then the acid distills as a constant-boiling mixture or azeotrope at 108.6oC.

The market for hydrochloric acid itself is limited, and efforts were made to convert it to other useful materials. Some success was achieved by Deacon in conversion to chlorine, but the most successful early process used for the production of chlorine was the Weldon process developed in 1870.

Example. The Weldon process is an oxidation-reduction process based upon the oxidizing properties of the natural mineral manganese dioxide. Manganese dioxide reacts with hydrochloric acid as shown below:

4HCl + MnO2 rarrow.gif (63 bytes) MnCl2 + 2H2O + Cl2

There is regeneration of the manganese dioxide by adding lime and limestone and blowing air through it:

2MnCl2 + CaO + CaCO3 + O2 rarrow.gif (63 bytes) 2CaCl2 + 2MnO2 + CO2

The Weldon process is a much more economic process than an older nonregenerative process. The production of chlorine by treatment of a mixture of salt and manganese dioxide with sulfuric acid had previously been in use:

4NaCl + 2H2SO4 + MnO2 rarrow.gif (63 bytes) 2Na2SO4 + 2H2O + MnCl2 + Cl2

The chlorine produced by either process could be absorbed in alkali water (Javelle water); alternatively, a liquid bleach could be prepared from a sludge of Ca(OH)2 and chlorine. A dry bleaching powder could be prepared by reaction of wet chlorine with CaO; the reaction is

CaO + Cl2 rarrow.gif (63 bytes) Ca(OCl)2.

The dry bleaching powder was in commercial production by 1799 and its production has continued to the present day. The effect of the Weldon process was to provide a much cheaper source of chlorine.

Modern industrial production of chlorine employs the even less costly electrolytic chloralkali process, of which details are given elsewhere.

Natural and synthetic soda have competed for markets since the days of Leblanc. In Europe, the market is dominated by the synthetic material, while in North America more than half of the soda is natural trona; the primary source is Wyoming, U.S.A. Half of the modern market for soda is the production of glass, and another quarter is used for other chemical manufactures.

The Leblanc process was superseded by the ammonia-soda Solvay process, which began significant commercial production in the 1870's and is still in use. Although the process has been known since 1811, its commercial development was both long and difficult due to the many practical difficulties rather than to strictly chemical ones.

Example. The chemistry of the Solvay process is based on the reaction of salt, ammonia, and carbon dioxide. If a concentrated solution of salt is saturated with ammonia and carbon dioxide is then added, NaHCO3 precipitates:

NaCl(aq) + NH3(g) + CO2(g) + H2O rarrow.gif (63 bytes) NH4Cl(aq) + NaHCO3

By heating the sodium hydrogen carbonate is converted to soda:

2NaHCO3 rarrow.gif (63 bytes) Na2CO3(s) + CO2 + H2O

The resulting carbon dioxide is recycled. The aqueous solution of NH4Cl can also be recycled by treatment with lime or slaked lime, which regenerates the ammonia:

2NH4Cl(aq) + CaO rarrow.gif (63 bytes) CaCl2(aq) + H2O + 2NH3

Items such as (N2, H2, C, NaCl, H2O, CaO) are all raw materials whose inputs to the process are continuous; the hydrogen is produced from natural gas or as a by-product of oil refining. All other materials are recycled, except for the products which are underlined.

Electrolytic Chloralkali Production

The alkali industry also includes electrolytic chloralkali production because electrolysis of concentrated aqueous solutions of sodium chloride produces sodium hydroxide as well as chlorine. This process is discussed in a different section. Much of the caustic soda (NaOH) and nearly all of the chlorine used in industry are produced in this way.

Upon electrolysis aqueous sodium chloride solutions liberate chlorine gas at the carbon anode where chloride ion is oxidized to chlorine; the reaction is

2Cl- rarrow.gif (63 bytes) 2e- + Cl2(g).

The mercury or carbon cathode is the site of reduction of water to hydrogen with concomitant production of hydroxide; the reaction is

2H2O + 2e- rarrow.gif (63 bytes) 2OH-(aq) + H2(g)

The sodium hydroxide solution remaining after the hydrogen and chlorine are removed is often evaporated to give solid sodium hydroxide; alternatively, the aqueous solution itself may be used for its alkali content.

Characteristics and Uses of Bases

Sodium hydroxide is the most commonly used strong base both in industry and in the laboratory for reasons of cost; potassium hydroxide is very similar in its behavior. Both are white waxy solids which are hygroscopic, attracting water vapor from the air to form concentrated aqueous solutions, and advantage can be taken of this in their use as laboratory drying agents. Concentrated aqueous solutions of both NaOH and KOH also remove carbon dioxide from air by the acid-base reaction:

2OH-(aq) + CO2(g) rarrow.gif (63 bytes) CO32-(aq) + H2O

As a consequence, solutions of these strong bases exposed to air react so as to reduce their hydroxide concentration. Such solutions are best stored in closed polyethylene bottles, since concentrated NaOH and KOH solutions also slowly react to etch glass.

Concentrated aqueous NaOH and KOH solutions are colorless and somewhat viscous. The usual laboratory concentrated sodium hydroxide solution is 50% NaOH by mass; it has a density of 1.53 kg/L which corresponds to a 19.1 molar aqueous solution. Concentrated aqueous solutions of strong bases are highly corrosive to the skin and severely damaging to the eye, and so must be handled with care.

When only the amount of base present is important, as in many industrial processes, NaOH and

Na2CO3 can be in direct competition and their prices often dictate the choice of alkali used. In North America, the major uses of NaOH are in the production of organic chemicals (35%), pulp and paper (20%), and inorganic chemicals (15%); soap and detergents now account for only 5% of the NaOH used.

Copyright 1997 James R. Fromm