Chemistry of Sulfur
Sulfur is a chemical element that is represented with the chemical symbol "S" and the atomic number 16 on the periodic table. Because it is 0.0384% of the Earth's crust, sulfur is the seventeenth most abundant element following strontium. Sulfur also takes on many forms, which include elemental sulfur, organo-sulfur compounds in oil and coal, H2S(g) in natural gas, and mineral sulfides and sulfates. This element is extracted by using the Frasch process, a method where superheated water and compressed air is used to draw liquid sulfur to the surface. Offshore sites, Texas, and Louisiana are the primary sites that yield extensive amounts of elemental sulfur. However, elemental sulfur can also be produced by reducing H2S, commonly found in oil and natural gas. For the most part though, sulfur is used to produce SO2(g) and H2SO4.
Sulfur at a Glance
Physical Properties of Sulfur
Sulfur has an atomic weight of 32.066 grams per mole and is part of group 16, the oxygen family. It is a nonmetal and has a specific heat of 0.706 J g-1 oC-1. The electron affinitry if 200 kJ mol-1 and the electronegativity is 2.58 (unitless).
Sulfur is typically found as a light-yellow, oqaque, and brittle solid in large amounts of small orthorhombic crystals. Not only does sulfur have twice the density of water, it is also insoluble in water. On the other hand, sulfur is highly soluble in carbon disulfide and slightly soluble in many common solvents. Sulfur can also vary in color and blackens upon boiling due to carbonaceous impurities. Even as little as 0.05% of carbonaceous matter darkens sulfur significantly.
Allotropy and Polymorphism of Sulfur
Compared to other elements, sulfur has the most allotropes. While the S8 ring is the most common allotrope, there are 6 other structures with up to 20 sulfur atoms per ring.
Oxides of Sulfur There are many different stable sulfur oxides but the two that are commonly found are sulfur dioxide and sulfur trioxide.
Sulfur dioxide is a commonly found oxide of sulfur. It is a colorless, pungent, and nonflammable gas. It has a density of 2.8 kg/m3 and a melting point of -72.5 oC. Because organic materials are more soluble in SO2 than in water, the liquid form is a good solvent. SO2 is primarily used to produce SO3. The direct combustion of sulfur and the roasting of metal sulfides yield SO2 (called the contact process):
S(s) + O2(g) → SO2(g) (Direct combustion)
2 ZnS(s) + 3 O2(g) → 2 ZnO(s) + 2 SO2(g) (Roasting of metal sulfides)
Sulfur trioxide is another one of the commonly found oxides of sulfur. It is a colorless liquid with a melting point of 16.9 oC and a density of kg/m3. SO3 is used to produce sulfuric acid. SO2 is used in the synthesis of SO3:
2 SO2(g) + O2(g) 2 SO3(g) (Exothermic, reversible reaction)
This reaction needs a catalyst to be completed in a reasonable amount of time. V2O5 is the catalyst most commonly used.
Sulfuric acid, H2SO4, is produced by reacting SO3 with water. However, this often leads to pollution problems. SO3(g) is reacted with 98% H2SO4 in towers full of ceramic material to produce H2S2O7 or oleum. Water is circulated in the tower to maintain the correct concentration and the acid is diluted with water at the end in order to produce the correct concentration. Pure sulfuric acid has no color and odor, and it is an oily, hygroscopic liquid. However, sulfuric acid vapor produces heavy, white smoke and a suffocating odor.
Dilute sulfuric acid, H2SO4(aq), reacts with metals and acts as a strong acid in common chemical reactions. It is used to produce H2(g) and liberate CO2(g) and can neutralize strong bases.
Concentrated sulfuric acid, H2SO4(concd), has a strong affinity for water. In some cases, it removes H and O atoms. Concentrated sulfuric acid is also a good oxidizing agent and reacts with some metals.
C12H22O11(s) → 12 C(s) + 11 H2O(l)
(Concentrated sulfuric acid used in forward reaction to remove H and O atoms.)
Sulfurous acid (H2SO3) is produced when SO2(g) reacts with water. It cannot be isolated in its pure form, however, it forms salts as sulfites. Sulfites can act as both reducing agents and oxidizing agents.
O2(g) + 2 SO32-(aq) → 2 SO42- (aq) (Reducing agent)
2 H2S(g) + 2 H+(aq) + SO32-(aq) → 3 H2O(l) + 3 S(s) (Oxidizing agent)
H2SO3 is a diprotic acid that acts as a weak acid in both steps and H2SO4 is also a diprotic acid but acts as a stong acid in the first step and a weak acid in the second step. Acids like NaHSO3 and NaHSO4 are called acid salts because they are the product of the first step of these diprotic acids.
Boiling elemental sulfur in a solution of sodium sulfite yields thiosulfate. Not only are thiosulfates important in photographic processing, but they are also common analytical reagents used with iodine (like in the following two reactions).
2 Cu2+(aq) + 5 I-(aq) → 2 CuI(s) + I3-(aq)
I3-(aq) + 2 S2O32-(aq) → 3 I-(aq) + S4O62-(aq) [Excess triiodide ion titrated with Na2S2O3(aq)]
Sulfur halides are compounds formed between sulfur and the halogens. Common compounds include SF2, S2F2, SF4, and SF6. While SF4 is a powerful fluorinating agent, SF6 is a colorless, odorless, unreactive gas. Compounds formed by sulfur and chloride include S2Cl2, SCl4, and SCl2. SCl2 is a red bad smelling liquid that is utilized to produce mustard gas.
SCl2 + 2CH2CH2 → S(CH2CH2Cl)2 (Mustard gas)
Sulfur has many practical applications. As a fungicide, sulfur is used to counteract apple scab in organically farmed apple production. Other crops that utilize sulfur fungicides include grapes, strawberries, and many vegetables. In general, sulfur is effective against mildew diseases and black spot. Sulfur can also be used as an organic insecticide. Sulfites are frequently used to bleach paper and preserve dried fruit.
The vulcanization of rubber includes the use of sulfur as well. Cellophane and rayon are produced with carbon disulfide, a product of sulfur and methane. Sulfur compounds can also be found in detergents, acne treatments, and agrichemicals. Magnesium sulfate (epsom salt) has many uses, ranging from bath additives to exfoliants. Sulfur is being increasingly used as a fertilizer as well. Because standard sulfur is hydrophobic, it is covered with a surfactant by bacteria before oxidation can occur. Sulfur is therefore a slow-release fertilizer. Lastly, sulfur functions as a light-generating medium in sulfur lamps.
Concentrated sulfuric acid was once one of the most produced chemicals in the United States, the majority of the H2SO4 that is now produced is used in fertilizer. It is also used in oil refining, production of titanium dioxide, in emergency power supplies and car batteries. The mineral gypsum is calcium sulfate dihydrate is used in making plaster of Paris. Over one million tons of aluminum sulfate is produced each year in the United States by reacting H2SO4 and Al2O3. This compound is important in water purification. Copper sulfate is used in electroplating. Sulfites are used in the paper making industry because they produce a substance that coats the cellulose in the word and frees the fibers of the wood for treatment.
Emissions and the Environment
Particles, SO2(g), and H2SO4 mist are the components of industrial smog. Because power plants burn coal or high-sulfur fuel oils, SO2(g) is released into the air. When catalyzed on the surfaces of airborne particles, SO2 can be oxidized to SO3. A reaction with NO2 works as well as shown in the following reaction:
SO2(g) + NO2(g) --> SO3(g) + NO(g)
H2SO4 mist is then produced after SO3 reacts with water vapor in the air. If H2SO4 reacts with airborne NH3, (NH4)2SO4 is produced. When SO2(g) and H2SO4 reach levels that exceed 0.10 ppm, they are potentially harmful. By removing sulfur from fuels and controlling emissions, acid rain and industrial smog can be kept under control. Processes like the fluidized bed combustion have been presented to remove SO2 from smokestack gases.
Note: Highlight the spaces for the answers. (Ctrl+A or use the mouse.)
1. Draw a diagram that summarizes the allotropy of sulfur. Use symbols, arrows, and numbers.
Answer: The diagram may be drawn in any way. However, the symbols (S?), (S?), (S?), (S?), and S8(g) must be included. The temperatures should be written next to the arrows.
2. Direct combustion of sulfur is the only method for producing SO2(g). True or False.
3. Sulfites are not oxidizing agents. They are good reducing agents. True or False.
4. Give the reaction for the production of sulfur trioxide.
Answer: 2 SO2(g) + O2(g) 2 SO3(g)
5. Choose the incorrect statement.
A. Sulfur produces cellophane and rayon.
B. Standard sulfur is hydrophobic.
C. SO2 can oxidize to SO3
D. Sulfur influences the development of acid rain and industrial smog.
E. All of the above are correct.
This page viewed 8989 times