Chemical & Chemical Engineering News (80th Anniversary Issue), Vol. 81, No. 36, 2003, Sept. Edited by X. Lu Introduction

But as a chemist turned playwright, let me end with some lines from “Oxygen”—a play I wrote with Roald Hoffmann:

ASTRID: First to the discovery: No one will question that oxygen confers great benefit on mankind, right?

BENGT: Oxygen was good for people before it was “discovered!”

And then Mme. Lavoisier’s conclusion of the play: “Imagine what it means to understand what gives a leaf its color! And how it turns red. What makes a fever fall, a flame burn. Imagine!”

Carl Djerassi is a playwright, novelist, and professor of chemistry emeritus at Stanford University. He has won both the National Medal of Science (for the first synthesis of a steroid oral contraceptive) and the National Medal of Technology (for promoting new approaches to insect control).

OXYGEN AT A GLANCE

Name: From the Greek oxy genes, acid forming. The name came from an incorrect belief that oxygen was needed to form all acids.

Atomic mass: 15.99.

History: The discovery of oxygen is usually credited to English chemist Joseph Priestley in 1774. It was discovered independently by Carl W. Scheele in Uppsala, Sweden, but published later.

Occurrence: O2 makes up about 20.95% of the atmosphere by volume. Ozone, O3, is a reactive gas; in the upper atmosphere, it blocks harmful solar radiation.

Appearance: Colorless, odorless gas at room temperature; pale blue as a liquid and a solid; faintly blue with a brackish odor as gaseous ozone.

Behavior: Oxygen supports combustion and combines with most elements to give both solid and gaseous oxides.

Uses: Essential for respiration. Oxygen is used in steelmaking, in metal cutting, and in the chemicals industry to make methanol and ethylene oxide.

SULFUR

AMOS B. SMITH III, UNIVERSITY OF PENNSYLVANIA

To many people, the element sulfur quickly brings to mind unpleasant memories of rotten eggs and unwelcome encounters with skunks. While true that many sulfur-containing compounds, such as mercaptans, thioethers, and disulfides, possess remarkably foul odors, elemental sulfur (S) in pure form is completely odorless, tasteless, and by and large nontoxic.

Referred to in Genesis as brimstone, meaning "a stone that burns," elemental sulfur, which does readily burn in air, has been known since ancient times. The word derives from sulvere and sulphur (Sanskrit and Latin, respectively). Major sulfur deposits are found in the salt domes ubiquitous to the Gulf Coast of the U.S.; sulfur is the 16th most abundant element in nature. Recovery is by melting with superheated water, a commercial process developed by Herman Frasch in the 1890s and still in use today.

Sulfur is also a major by-product of and can be extracted from coal, ores, and minerals such as gypsum, cinnabar, barite, and pyrite (fool's gold). From an economic perspective, most of the world's sulfur production is used to make sulfuric acid (H₂SO₄), approximately 40 million tons every year just in the U.S. Fertilizers and lead-storage automobile batteries consume a large portion of this supply, with smaller amounts used as insecticides, as dyeing agents, in the manufacture of gunpowder, and to vulcanize natural and synthetic rubbers to impart desirable mechanical properties.

Since sulfur is positioned directly under oxygen in the periodic table, one might suspect relatively simple behavior at the atomic level. This is certainly not the case! Indeed, Jerry Donohue, in his authoritative treatise, "The Structures of the Elements," argues, "Of all of the elements, sulfur presents the most confusion and complexity in this respect." Unlike sulfur's near neighbors, such as N, O, Cl, and Br, or P, As, Se, and Te, which in elemental form exist respectively as diatomic molecules or take on polymorphic networks, the four isotopes of sulfur [³²S (95.1%), ³³S (0.74%), ³⁴S (4.2%), and ³⁶S (0.016%)] combine to form a uniquely large number of allotropic forms (compare S_n, where n = 1–12, 18, and infinity). Of these, only eight have been characterized crystallographically. The most common form, stable at room temperature and atmospheric pressure, is orthorhombic sulfur (S₈). Here, eight sulfur atoms bond covalently in crownlike rings. This allotrope, also known as rhombic sulfur, Muthmann's sulfur, and a-S, was among the first substances to be examined crystallographically by William Bragg in 1914.

The chemistry of sulfur is equally complex; it combines with many elements to yield a bewildering array of both organic and inorganic compounds. Common inorganic compounds include sulfur hydrides, carbon disulfide (CS2), sulfur selenides, and sulfur halides (SX6). Oxides of sulfur are particularly important, possessing both beneficial and deleterious properties. Sulfur dioxide (SO2), for example, finds beneficial use in preserving fruits and vegetables and in the brewing and wine-making industry as both an antioxidant and an antibiotic. Sulfur dioxide and its close relative, sulfur trioxide (SO3), represent serious hazards, arising in the environment principally by burning of sulfur-rich fuels such as coal and oil or by smelting ores. Released into the atmosphere and combined with water, these pollutants form sulfuric acid and in turn acid rain, a cause of huge economic damage.

	CLUMPED Sulfur produces a monoclinic shape when crystallized out of solution.