The Acid Industry

James Richard Fromm

Sulfuric Acid

The fundamental compound of the acid industry has always been sulfuric acid, H2SO4. Sulfuric acid has been known since the days of alchemy under the name of oil of vitriol due to its oily and viscous appearance. Its preparation, known from 1570 onwards, was by distillation of the natural mineral "green vitriol", FeSO4.7H2O:

FeSO4.7H2O rarrow.gif (63 bytes) FeO + H2SO4 + 6H2O

The dependence of the process upon green vitriol, which is not a common mineral, and low yield led to its replacement by the per campanem (by the bell) method prior to 1700. In this method, sulfur and potassium nitrate (saltpeter; KNO3) were burned together under large glass bell-shaped containers. This process was used until 1746, when it was superseded by the lead chamber process The lead chamber process, chemically identical to the per campanem process, permitted large-scale sulfuric acid production. In both methods, the reaction of sulfur dioxide with oxygen requires the oxides of nitrogen as catalysts. The steps are:

S + O2 rarrow.gif (63 bytes) SO2(g)

2SO2 + O2 rarrow.gif (63 bytes) 2SO3(g)

SO3 + H2O(l) rarrow.gif (63 bytes) H2SO4

The net overall reaction is:

2S + 3O2 + 2H2O rarrow.gif (63 bytes) 2H2SO4

In the eighteenth century sulfuric acid was in use in pharmacy for manufacture of Na2SO4 and for the removal of metals and metal oxide films by metal workers. The production of bleach demanded more of the acid, and demand was greatly increased by the rise of the synthetic soda industry. Although other production methods may now be more economical for some of them, all of the other strong acids can be, and have been, commercially produced from sulfuric acid:

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

2KNO3 + H2SO4 rarrow.gif (63 bytes) 2HNO3 + K2SO4

Pyrite (FeS), which on roasting yields SO2, is often used to replace sulfur for production of sulfuric acid. Much of the sulfuric acid of commerce is a by-product of metal production, employing the SO2 evolved in the roasting of sulfide metal ores. Sulfuric acid producers use the cheapest source of raw materials. The shift in England from sulfur to pyrite in 1840 was the result of a tripling of the price of sulfur exported from Sicily. More recent shifts between sulfur by-products from natural gas desulfurization and metallurgical processing have been governed by economic pressures imposed by antipollution legislation.

The oxides of nitrogen serve as a catalyst in the production of SO3 from SO2. Originally they were exhausted as waste, but their recovery became economically significant after 1860. Other catalysts, such as finely divided platinum, are also effective. Platinum can be used to produce oleum (H2SO4 containing excess SO3, in demand for synthetic organic chemical production since 1870) which cannot be made by the lead chamber process. Both are still in use, the catalytic or contact process probably now being more common.

Sulfuric acid in concentrated form is a common laboratory reagent. Reagent-grade sulfuric acid is a colorless oily liquid; it is about 94% H2SO4, the remainder being water. It has a density of 1.83 kg/L which corresponds to an aqueous solution of 17.6 molar H2SO4. Concentrated sulfuric acid has a very strong affinity for water and reacts to rapidly dehydrate and attack both flesh and clothing, so it must be handled with care. There is a large heat of reaction for the reaction between concentrated sulfuric acid and water which is sufficient to boil a small amount of water added to concentrated sulfuric acid, ejecting acid from the container.

When concentrated sulfuric acid and water must be mixed, it is far safer to add the acid to the water rather than the water to the acid. Even so, this dilution must be done with care.

Most of the sulfuric acid produced in North America, about 70%, is used in the manufacture of fertilizers. Metals recovery, petroleum refining, and chemical manufacturing each account for another 5%; the remaining 15% has many smaller uses.

Hydrochloric Acid

Hydrochloric acid is gaseous HCl or its aqueous solutions, since HCl is highly soluble in water. It dissociates to H3O+(aq) and Cl-(aq) virtually completely and is one of the common strong acids. Hydrogen chloride gas can be prepared easily from common salt and sulfuric acid, the reactions being:

NaCl(c) + H2SO4(l) rarrow.gif (63 bytes) NaHSO4(c) + HCl(g) and

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

Both reactions are used industrially, the reaction to NaHSO4 predominating at 150oC while the reaction to Na2SO4 predominates at 550oC. The Leblanc process, an early source of HCl, has now been replaced by production from salt and sulfuric acid. Other production methods for HCl include the burning of chlorine, produced by the electrolytic chloralkali process, in hydrogen:

H2(g) + Cl2(g) rarrow.gif (63 bytes) 2HCl(g)

This process is particularly useful because hydrogen is produced in the same chloralkali process that produces chlorine, so reaction of the two gases coming from the same cell yields HCl with no by-products. A chlorine burner using this reaction can produce a high-concentration, high-quality product. However, the reaction can proceed explosively and so careful handling and appropriate safety devices are necessary.

The concentrated hydrochloric acid of commerce is normally the azeotrope. Impure hydrochloric acid is sometimes yellow and is known by the older, if not archaic, name of muriatic acid, but concentrated reagent-grade hydrochloric acid is colorless. Its density is 1.19 kg/L, which corresponds to a 12.4 molar aqueous solution of HCl.

Copyright 1997 James R. Fromm