Volume 3

Natural Occurrence and Physical Properties • Sulfur Oxides. Sulfurous Acid and Sulfuric Acid. The Contact Process • Hydrogen Sulfide • Exercises
3.4.1. Natural Occurrence and Physical Properties. Sulfur is one of the elements of Group 16 of the periodic table. These elements are often called the chalcogens. Sulfur is the 17th most abundant element in the Earth's crust, occurring both as a simple substance and in the form of various minerals (Figure 3-33). The most common mineral of sulfur is pyrite, FeS2, also known as fool's gold (Figure 3-33, right). More than 30 allotropes of sulfur are known. The most stable allotropic form is orthorhombic sulfur, a soft crystalline yellow solid (Figure 3-33, left), consisting of cyclic molecules formed by 8 sulfur atoms each. It is for this reason that the formula of sulfur is sometimes written as S8 in chemical equations. In our course, we will use just "S" for reactions of sulfur as a simple substance.

Pure elemental sulfur is odorless. The unpleasant "sulfur smell" is that of many sulfur compounds, not sulfur itself. The melting and boiling points of sulfur are 120 oC and 445 oC, respectively.
Figure 3-33. Samples of native sulfur (left; source) and pyrite, FeS2, also known as fool's gold (right; source).

3.4.2. Sulfur Oxides. Sulfurous Acid and Sulfuric Acid. The Contact Process. Sulfur forms two main oxides, sulfur dioxide, SO2, and sulfur trioxide, SO3. Burning sulfur in air gives rise exculsively to SO2, a colorless toxic gas with a pungent smell. To make SO3, SO2 is oxidized with the oxygen of the air in the presence of a catalyst, such as V2O5 (Figure 3-34). A toxic compound and an irritant, SO3 melts at 17 oC and boils at 45 oC. Also, two polymer forms of SO3 are known, one melting at 62 oC (α-SO3) and the other at 33 oC (β-SO3).
Figure 3-34. Reaction of sulfur with oxygen (left) and catalytic oxidation of SO2 to SO3 (right).

Both SO2 and SO3 are acidic oxides that react with water to give H2SO3 (sulfurous acid) and H2SO4 (sulfuric acid), respectively (Figure 3-35).
Figure 3-35. Reactions of SO2 and SO3 with water.

The reaction of SO2 with water is reversible. As a result, solutions of H2SO3 smell SO2. In contrast, SO3 reacts with water to give odorless H2SO4 irreversibly. Unlike sulfuric acid, H2SO3 is a weak acid.

The oxidation state of sulfur in SO2 as well as in H2SO3 and its salts is +4. All of these S (+4) compounds are reductants that are easily oxidized to sulfate where the S atom has the oxidation state of +6. Two chemical equations presented in Figure 3-36 are for the reduction of iron (+3) to iron (+2) and of iodine (0) to iodine (-1) with SO2 or H2SO3.
Figure 3-36. Two examples of redox reactions involving reduction by SO2 or H2SO3.

Sulfuric acid is the number one chemical produced worldwide. The process to make H2SO4 involves the synthesis of SO2, followed by its catalytic oxidation to SO3. Although some SO2 for the manufacturing of H2SO4 is obtained by combining naturally occurring sulfur and oxygen (air), larger quantities of SO2 are made by burning the mineral pyrite, FeS2 (Figure 3-33, right), the most abundant natural source of sulfur. In industry, this process is sometimes called the weathering of pyrite. The SO2 gas produced is then purified, mixed with extra air, and passed over a V2O5 catalyst at approximately 450 oC to bring about the oxidation to SO3 (Figure 3-37). This two-step process is called the contact sulfuric acid process or just the contact process. The name comes from the key step, in which a gaseous mixture of SO2 and air is put in contact with the solid vanadium catalyst.
Figure 3-37. The contact process for making H2SO4.

To make H2SO4 from SO3, one just needs to react SO3 with water (Figure 3-35). However, the industrially produced SO3 is never treated with water because the reaction between the two is very exothermic, producing a highly corrosive mist that takes forever to settle. Instead, the SO3 from the contact process is absorbed into already available concentrated H2SO4 to produce oleum, a solution of SO3 in H2SO4 (Figure 3-37). To make concentrated sufluric acid, oleum is mixed with a calculated amount of water.

3.4.3. Hydrogen Sulfide. Hydrogen sulfide, H2S, is a toxic gas that is notorious for its characteristic rotten egg odor. Not only is H2S toxic, but it is also insidiously toxic. After inhaling H2S for minutes, an individual loses his sensitivity to the odor (olfactory fatigue) and might think that (s)he is no longer exposed to this poisonous gas. Breathing high levels of hydrogen sulfide can be fatal.

As already mentioned earlier in our course, H2S is a very weak acid. A standard laboratory method to make H2S is the treatment of FeS with hydrochloric acid (Figure 3-38).
Figure 3-38. Reaction of FeS with HCl is a laboratory method to make H2S.

In H2S, the sulfur atom has the oxidation state of -2, the lowest possible oxidation state for sulfur. Therefore, H2S is devoid of oxidizing ability and can only serve as a reductant. An interesting comproportionation reaction between H2S and SO2 gives sulfur in the form of a fine powder (Figure 3-39).
Figure 3-39. Comproportionation of H2S and SO2 to elemental sulfur.

Hydrogen sulfide finds applications in analytical chemistry, paper making, organic synthesis, and separation of heavy water (deuterium oxide, D2O) from natural water.

3.4.4. Exercises.

1. Write the electronic configuration of a sulfur atom and account for the oxidation state of -2 of S in H2S. Explain why H2S cannot play the role of an oxidizer. Answer

2. What sulfur mineral is called fool's gold? What are the formula and main application of this mineral? What are the oxidation states of the elements in fool's gold? Answer
3. Sulfur burns in air in the absence of a catalyst to give (a) exclusively SO2; (b) exclusively SO3; (c) a mixture of SO2 and SO3; (d) sulfur does not burn in the absence of a catalyst. Answer

4. Write balanced equations for the reactions between (a) S and O2; (b) SO2 and O2 (in the presence of a vanadium catalyst); (c) SO2 and H2S; (d) I2 and H2SO3 (SO2 + H2O); (e) FeS and HCl; (f) H2SO3 and MnO2 (products: MnSO4 and H2O); (g) SO3 and H2O; (h) SO2 and NaOH (excess); (i) SO3 + Ca(OH)2; (j) H2S and KOH (excess). Find redox reactions and identify the oxidant and the reductant. Indicate changes in the oxidation states. Answer

5. What is the contact process for making H2SO4? Where does the name "contact process" come from? Write chemical equations for the reactions employed in the contact process. [Answer: See section 3.4.2]

6. What is oleum? Why is the SO3 produced in the contact process dissolved in already existing H2SO4 rather than reacted with water to make H2SO4? How is H2SO4 made from oleum? [Answer: See section 3.4.2]

7. Both SO2 and SO3 have a pungent irritating odor. Both react with water to give the corresponding acid (sulfurous and sulfuric, respectively). However, concentrated aqueous solutions of sulfurous acid still smell SO2, whereas solutions of H2SO4 are odorless. Why? Answer

8. How much water should be used for mixing with 100 g of 30% oleum to convert it to 100% H2SO4? Answer

9. Sulforous acid H2SO3 and hydrogen sulfide H2S are weak acids. True or false? Answer

10. The pH of an aqueous solution of K2SO3 is (a) approximately 7; (b) >7; (c) <7. Answer

11. A student criticized the upper equation in Figure 3-36.

I2 + SO2 + 2 H2O = 2 HI + H2SO4

He insisted that since SO2 reacts with water to give H2SO3, the equation should be written as follows

I2 + H2SO3 + H2O = 2 HI + H2SO4

Which of the two equations is correct? Answer