18.2.

Remember from Chapter 7 that water (H2O) can be viewed as hydrogen hydroxide (HOH) and some small fraction (1 in 10,000,000) will ionize to form hydrogen (H+) and hydroxide (OH-) ions:

In pure water, the concentrations of hydrogen and hydroxide ions are equal. Anything that causes the concentration of hydrogen ion to exceed that of hydroxide ion is called an acid; anything that causes the concentration of hydroxide ion to exceed that of hydrogen ion is called an alkali, or base. Now, we have spent a lot of time in this book talking about alkalis: potash and soda ash in Chapter 8, lime in Chapter 10, ammonia in Chapter 12, and Caustic Soda in Chapter 15. Alkalis have a bitter taste, a high pH (8-14), and turn pH test paper blue. This chapter is an introduction to acids.

The earliest, and for most of human history the only acid available was vinegar, a solution of acetic acid in water produced by the bacterial oxidation of alcoholic beverages as shown in Equation 4-1(c). Like vinegar, acids in general have a sour taste, a low pH (0-6), and turn pH test paper red. Acetic acid dissociates directly to produce hydrogen ion in aqueous solution, but there are other compounds, particularly non-metal oxides, that combine with water before dissociating. One such oxide is carbon dioxide.

Remember from Chapter 10 that when carbon dioxide dissolves in water, it forms hydrogen carbonate, AKA carbonic acid (H2CO3). Carbonic acid ionizes in water to form a bicarbonate[1] ion (HCO3-) and a hydrogen ion. If you add a base, say, sodium hydroxide to carbonic acid you get a salt, sodium bicarbonate. The hydroxide ion from the base combines with the hydrogen ion from the acid to produce water; the sodium ion left over from the sodium hydroxide and the bicarbonate ion left over from the carbonic acid stay in solution. You would have exactly the same situation if you just dissolved sodium bicarbonate in water and so in the equation, we abbreviate "Na+(aq) + HCO3-(aq)" as "NaHCO3(aq)." The bicarbonate ion is called the conjugate base of carbonic acid. You could also say that carbonic acid is the conjugate acid of the bicarbonate ion. The meaning of the word conjugate here is just that carbonic acid and bicarbonate ion are really the same thing, with and without an extra hydrogen ion.

We can take it even farther by adding sodium hydroxide to sodium bicarbonate. The hydroxide ion pulls a hydrogen ion off of the bicarbonate ion, making water and leaving carbonate ion (CO32-). The sodium ion from the sodium hydroxide joins the one from the sodium bicarbonate, so now you have two of them in solution. Same as before, you can abbreviate "2 Na+(aq) + CO32-(aq)" as "Na2CO3(aq)." Carbonate ion is the conjugate base of bicarbonate ion and bicarbonate ion is the conjugate acid of carbonate ion. The acid-base chemistry of carbon dioxide is summarized in Equation 18-1; the oxides of sulfur behave similarly and are the main focus of this chapter.

Equation 18-1. Acid Properties of Carbon Dioxide

There are two oxides of sulfur: sulfur dioxide (SO2) and sulfur trioxide (SO3). When sulfur burns in air the product is sulfur dioxide, which we saw in Equation 9-2 as one of the products of the roasting of metal sulfide ores. Sulfur dioxide is a versatile compound; in the presence of a strong oxidant it plays the role of a reducing agent; in the presence of a strong reductant it acts as an oxidizing agent. As shown in Equation 18-2(b), it dissolves in water to form an acid, in this case a weak acid called sulfurous acid (H2SO3), which smells of rotten eggs. A mole of sulfurous acid reacts with a mole or two of sodium hydroxide to produce the salts, sodium bisulfite and sodium sulfite, respectively. While there has never been a large market for sulfurous acid or its salts, they are not entirely useless, as we'll see in Chapter 24. But by far the biggest demand in the acid world has always been for sulfuric acid.

Equation 18-2. Properties of Sulfurous Acid

The production of sulfuric acid (H2SO4) requires a bit of ingenuity. As we have seen, combustion of sulfur by atmospheric oxygen produced sulfur dioxide, not sulfur trioxide. In Chapter 17 we saw that saltpeter could substitute for atmospheric oxygen in a redox reaction. Ward's contribution was to prepare a modified "gunpowder," one consisting of only sulfur and sodium nitrate.[2] When such a mixture is burned, it produces sulfur dioxide and nitrogen oxide, as shown in Equation 18-3(a). The nitrogen oxide liberated from the saltpeter reacts with atmospheric oxygen to produce nitrogen dioxide, which, in turn, converts sulfur dioxide to sulfur trioxide. If you add the three reactions of Equation 18-3(a), the nitrogen dioxides cancel out:

Equation 18-3. Properties of Sulfuric Acid

3 S + 2 NaNO3 + O2 = Na2S + 2 SO3 + 2 NO

If more O2 were available, the NO would be converted back into NO2 and could react with more SO2 to produce more SO3. But that would regenerate NO, which could react with more O2 to make—well, you get the picture. The NO is not used up in this reaction; it simply changes to NO2 and back again. We say that nitrogen oxide is a catalyst for the oxidation of sulfur dioxide to sulfur trioxide. The entire process is summarized in schematic form in Figure 18-1. The first reactor is an old friend, a burner, first encountered in Figure 1-3. Equation (a) of Figure 18-1 is shorthand for Equation 18-3(a), the "NO" over the equal sign denoting the catalytic action of nitrogen oxide. The second reactor is called an absorber, which consists of nothing more than a container in which a gas can dissolve in a liquid.

Figure 18-1. The Lead Chamber Process

The sulfuric acid of commerce is 95% sulfuric acid and 5% water. You can buy it in this concentration, for example, as industrial-strength drain opener.[3] The acid used to fill car batteries is a solution of sulfuric acid in water. But in this chapter you will learn to make sulfuric acid from scratch.

Sulfuric acid reacts with a variety of substances. Given its name, we would expect it to react with bases to form salts. A mole of sulfuric acid reacts with a mole or two of sodium hydroxide, for example, to produce the salts, sodium bisulfate and sodium sulfate, respectively. Sulfuric acid can also oxidize metals, for example reacting with iron to produce iron sulfate and hydrogen gas. Finally, sulfuric acid is often used as a dehydrating agent. It will, for example, suck water right out of sugar, leaving charcoal: the same reaction we saw way back in Equation 1-1. You can think of sulfuric acid as a kind of chemical bejeezus sucker. It is this triple nature of sulfuric acid, as an acid, as an oxidizing agent, and as a dehydrating agent, which accounts its place as the king of chemicals.

Sulfuric acid is certainly the most important of the acids, if raw tonnage is a measure, but there are two other acids of commercial importance, and it turns out that sulfuric acid can be used to manufacture them both. One of these, nitric acid, HNO3, is an even more powerful oxidizing agent than sulfuric acid. It will oxidize even those metals like lead, copper, and silver which are resistant to sulfuric acid. It will turn out to be important in the manufacture of such amazing applications as photographic film, fertilizers, explosives, and plastics. Nitric acid can be distilled from saltpeter and sulfuric acid, as shown in Equation 18-4. The other acid, hydrochloric acid, HCl, will play a significant role in the development of the chemical industry in the nineteenth century. It can be distilled from ordinary table salt and sulfuric acid. This triumvirate of mineral acids has played such a decisive role in the development of the world as we know it that they deserve a little mnemonic to help you remember their importance:

 

In this way was the World created. From this there will be amazing applications because this is the pattern. Therefore am I called Hermes Trismegistos, having the three parts of wisdom of the Whole World.

 The Emerald Tablet of Hermes Trismegistos

Equation 18-4. Two More Mineral Acids

WarningMaterial Safety
 

Locate an MSDS's for saltpeter (CAS 7757-79-1), sulfur (CAS7704-34-9), and sulfuric acid (CAS 7664-93-9). Summarize the hazardous properties in your notebook, including the identity of the company which produced each MSDS and the NFPA diamond for each material.[4]

Your most likely exposure will be to sulfur dioxide fumes. If you are careless enough to get a lung-full and if it makes you cough more than once or twice, you should go to the hospital. But you would be better off not getting a lung-full; just keep your nose out of where the gas is.

You should wear safety glasses while working on this project. All activities should be performed either outdoors or in a fume hood. Leftover sulfur may be thrown in the trash. Leftover saltpeter and sulfuric acid can be flushed down the drain with plenty of water.

NoteResearch and Development
 

Before you get started, you should know this stuff.

Notes

[1]

The modern name for the bicarbonate ion is the hydrogen carbonate ion. The older name, however, continues to be widely used.

[2]

Sodium nitrate is less expensive and less explosive than potassium nitrate; a double benefit for the budding acid maker.

[3]

The solid drain openers available at grocery stores are almost universally caustic soda, sodium hydroxide. Industrial-strength liquid drain openers are more commonly sulfuric acid. Mixing the two is dangerous.

[4]

The NFPA diamond was introduced Section 15.2. You may substitute HMIS or Saf-T-Data ratings at your convenience.