Introduction to Chemical Species
The term species is rarely used in chemistry. Yet a figurative use of the term is appropriate because we often find a group of similar compounds, ions and elements that transform into one another as part of cycles, not only in nature but in industry. To acquaint ourselves with these species more intimately, while conveying their chemistry, we will not restrict ourselves to sights and sounds.
With chemistry it’s all to easy to do just that. It involves sizzling, refluxing sounds and explosions. Its stuff can be brilliantly colored and interactions with energy lead to dazzling changes like fireworks. Since electronic tools of communication emit electromagnetic waves and sound, which are forms of energy; since books and lectures transmit chemistry’s symbols and models through our eyes and ears, it’s easy to forget that our other senses can directly interact with matter. Given that chemistry is after all the study of matter, let’s remember to write about chemical species that can be tasted, felt and smelled, especially since the latter action is the best at evoking memories.
Sulfuric acid (H2SO4)
In concentrated form, sulfuric acid is a thick syrupy liquid. I recall using it while working in a quality control lab analysing the protein content of hot dogs. Despite the pressure differences of the fumehood which are meant to divert the acidic fumes away from the working environment, one of my colleagues became itchy all over her arms and torso. She reported to the company doctor, but regretted her decision when she felt that he had taken advantage of her sensitive reaction to the acidic fumes. He had asked her to disrobe and expose her breasts. It was the first time in my young life that I had heard of unprofessional behaviour from an esteemed individual.
Industry’s most widely produced compound, H2SO4, is used in lead batteries, copper refining and especially in the production of phosphoric acid for phosphate fertilizers. It’s also needed to make other acids, pharmaceuticals and dyes. Currently, sulfuric acid is made industrially by burning sulfur and by catalytically oxidising the subsequent product (SO2) with vanadium(V) oxide. Contrary to many internet sources, the new product, SO3, is not added to water in the final step. That reaction is too exothermic for a large-scale operation. Instead SO3 reacts with previously-made sulfuric acid to make H2S2O7, which then reacts with water to yield the acid.
When SO2 escapes into the environment, it becomes sulfuric acid through the reaction of hydroxyl radicals and moisture. This main component of acid rain has a serious impact on lungs, infrastructure and lakes. For this reason industries equipped with the latest technology have scrubbers. Some of these use limewater to convert the SO2 into calcium sulfite and eventually into the useful CaSO4 · 2H2O, which is the material in wallboard.
The sulfate ion (SO4 2-)
Sulfate is a part of the sulfuric acid molecule, but if it’s bonded to a metal, the compound will not share acid’s most distinguishing properties. The sulfate ion itself does not burn skin or taste bitter. SO4 2- is also part of the planet’s natural sulfur cycle. Algae and plants absorb sulfate ion to make a pair of amino acids that contain sulfur, cysteine and methionine.
Humans cannot make their own methionine. Obtaining it from our diets, we ultimately rely on other organisms’ ability to do something with sulfate. Even when we make cysteine we get the sulfur from methionine. The sulfur linkages between different cysteine molecules in a peptide chain help it fold and stabilise into a specific three dimensional protein that define its role. The importance of shape and role can range from how curly your hair is to the proper functioning of enzymes in your cells.
Without methionine no nucleated cell can even begin to make a protein. Bacteria in ruminants also use sulfate to produce cystine and methionine. This was discovered by using sulfate with a radioactively tagged isotope of sulfur. Elsewhere sulfate is reduced to H2S by specialized bacteria while others can oxidize either H2S or elemental sulfur back to sulfate.
Bear in mind that H2S is also produced by volcanoes along with sulfur dioxide, which can react in the atmosphere to yield sulfates again. Sulfate-containing particles are important in cloud formation; they serve as seeds for the condensation of water vapour.
Orthorhombic sulfur (S8) is the most common of several possible, so-called allotropes, molecular combinations involving a single element. It is one of the few elements known to the ancients because it can occur unbonded to other elements near volcanoes and hot springs. Egyptians used it on wheat 2000 years ago to prevent rust disease. The element is bright yellow and if you handle it and feel its smooth texture, your hands will stink. When I was young, a kid in our neighbourhood invited me to take a whiff of a sample of burning sulfur. (He had lit a sample from his chemistry set with a match.) It turned to an attractive blue colour. The colour incited me to move closer, but the smell felt like needles going up my nostrils. Of course the kid laughed at my reaction.
Although there are “sulfur-shy” plants such as gooseberries, apricots and raspberries that should never be treated with sulfur, others tolerate the element, which can prevent powdery mildew, rose black spot and fungal rust. Since sulfur is a preventive fungicide it will not work if the diseases have already appeared.
Sulfur dioxide (SO2)
SO2 is a useful gas in the production of sulfuric acid, but it is also an undesirable combustion product of coal, which is typically 1 – 4% sulfur. Nickel and copper ores are sulfur compounds, so in the roasting process, sulfur dioxide is again produced. Scrubbers are essential in converting the smelly and lung-irritating gas to useful products, and these were discussed in the sulfuric acid blurb.
Volcanoes are a natural source of sulfur dioxide. I once hiked near Kīlauea volcano in Hawaii and the SO2-containing volcanic smog was foul-smelling and made it difficult to breathe. During major eruptions, SO2 can be injected to an altitude of over 10 km and into the stratosphere. There, SO2 is converted to sulfate aerosols which reflect sunlight and have a temporary cooling effect on the Earth’s climate. They also deplete ozone.
The overall volcanic contribution of SO2 to our atmosphere, however, pales in comparison to the amount we discharge. Emissions in North America and Europe peaked in the mid-70s but are currently at 1/3 and ¼ of their maximum values. But meanwhile in east Asia, the emissions have reached Europe’s highest output from 40 years ago.
Sodium bisulfite (NaHSO3)
Sodium bisulfite, a source of the bisulfite ion, is made by reacting washing soda (sodium carbonate) with sulfur dioxide, the same starting point in the contact process-production of sulfuric acid. Water treatment plants use bisulfite to eliminate excess chlorine from both drinking water and treated sewage. Relative to chlorine, bisulfite acts as a reducing agent, allowing chlorine to gain electrons and become chloride. In a similar reaction, bisulfite can react with oxygen to prevent corrosion of large pipes. An oxidation-reduction reaction involving bisulfite has a strong, peculiar smell that you can feel in your throat because the bisulfite ion and an acidic proton are in equilibrium with water and sulfur dioxide. Amateur wine makers use bisulfite as a reducing agent to prevent the oxidation and discoloration of white wine. They have to be careful with concentrations; otherwise too much will leave a sulfurous taste that ruins the wine. Even when professionals are careful, they run up against individuals who are allergic to bisulfite, which makes them wheeze and develop hives.
Hydrogen sulfide (H2S)
The odour threshold for hydrogen sulfide gas is only 10 parts per billion to about 1.5 parts per million. It stinks like rotten eggs or like the nastiest type of flatulence, both of which involve the microbial breakdown of sulfurous amino acids into hydrogen sulfide. Luckily, we are sensitive to the smell because at 1 to 2 parts per thousand it causes instant death. In between the odor threshold and the lethal dose, it has unpleasant effects ranging from nausea, pulmonary edema and loss of consciousness. Since hydrogen sulfide is produced naturally from decaying organic matter, it can be released from sewage sludge and manure. Sulfur hot springs and natural gas are also sources. Some industrial processes that release it as a byproduct include wastewater treatment, petroleum refining and mining. Workers have died when ventilation and detection systems have been inadequate. Marriages have been threatened by the indiscretions of a partner suffering from H2S flatulence.
- Purdue University. Disease Management Strategies www.extension.purdue.edu/extmedia/bp/bp-69-w.pdf
- United States Department of Labor. Occupational Safety and Health Administration https://www.osha.gov/SLTC/hydrogensulfide/hazards.html