The Disappearing Spoon: And Other True Tales of Madness, Love, and the History of the World from the Periodic Table

Most people can recall a large chart on the wall of their high school chemistry class: the periodic table, filled with helpful information about the atomic elements but in many ways hard to grasp. Kean notes, “People remember the table with a mix of fascination, fondness, inadequacy, and loathing” (12). The chart looks like an uneven wall of square bricks, eighteen wide by 9 deep. Each brick represents an element, and each element’s location is crucial because “the coordinates of an element determine nearly everything scientifically interesting about it” (13).

The ancient philosopher Plato believed that everything longs for its missing half, as men and women long for each other. Indeed, atoms contain electrons that orbit in concentric shells, “like clouds swirling around an atom’s compact core” (18), and “each level needs a certain number of electrons to fill itself and feel satisfied” (16). Atoms will swap electrons with other atoms to fill those shells, in the process bonding to form chemicals.

One exception is helium, which contains two electrons that fill up its only shell: “[I]t doesn’t need to interact with other atoms or share or steal electrons to feel satisfied. Helium has found its erotic complement in itself” (16). Similarly, the gases on the right side of the periodic table—neon, argon, krypton, xenon, and radon— combine with no other atoms; their aloofness gives them their moniker, “the noble gases.” To Plato, these would be the rare perfect elements, “incorruptible and ideal” (17).

To the left of the noble gases on the table are the highly active halogens; at the far left are the alkali metals, “even more violent” elements (17). Alkalis have one more electron than needed to fill their shells; halogens have one fewer; the two types bond tightly to each other. A common example is sodium chloride, or table salt. The rest of the elements bond to each other with lesser degrees of enthusiasm. Calcium, with two extra electrons, will bond with two atoms of chlorine, each of which needs one electron to fill its outermost shell; the chemical that results is calcium chloride (CaCl2).

In the 1920s and 1930s, Nobel nominee Gilbert Lewis figured out that, in a reaction between an acid and a base, “an acid is an electron thief,” and bases such as bleach or lye “might be called electron donors” (21). Acids vary widely in strength: one antimony-based superacid is “100,000 billion billion billion times more potent than stomach acid and will eat through glass” (22). Antimony is found in ancient paints and mascara, and later in toxic medicines.

Kean goes on to state, “As we move horizontally across the periodic table, each element has one more electron than its neighbor to the left” (24). The way atoms store their electrons varies, and for those in the middle of the periodic table, the transition metals, things get complicated. These atoms store their electrons oddly, burying them in lower energy levels where they can’t interact with other atoms; thus, many of these atoms behave very similarly.

Two rows of atoms, numbered 57 through 71, lie somewhat out of order at the bottom of the chart; these are the lanthanides or rare earths. According to Kean, “[t]he lanthanides bury new electrons even more deeply than the transition metals, often two energy levels down” and “can barely be distinguished from one another” (26).

Next, Kean discusses the makeup of atoms. In the center is the nucleus, made of protons and neutrons. Protons have an electric charge opposite from electrons. The number of protons in the nucleus determines the type of atom. The number of neutrons can vary; these variations are called isotopes. Most of an atom’s mass is in its nucleus: “The atomic number plus the number of neutrons is called the atomic weight” (28).

The chapter concludes with a discussion of the elements’ complexity. Kean notes that hydrogen is both the simplest element in the universe and the most abundant and that helium is the second-simplest element and second most abundant (29). The third most abundant element, however, isn’t number three, lithium, but element eight, oxygen.

Dr. Maria Goeppert-Mayer overcame bias against women in science, obtained a university position, and, in the late 1940s, discovered that “protons and neutrons in the nucleus sit in shells just like electrons and that filling nuclear shells leads to stability” (29). Oxygen is doubly stable, both in the nucleus and in its electron shells, “which explains its seeming overabundance” (30). Goeppert-Mayer won the Nobel Prize for her work. 

Kean opens the chapter by describing amino acids, which contain large amounts of carbon, “the sixth (and most versatile) element on the periodic table” (34). Carbon has four electrons in its outer shell; it loans these out to a variety of other atoms when they bond, making carbon flexible and adaptive. Silicon is also flexible because, like carbon, it has an outer electron shell with four empty spaces (36).

Kean questions whether life could be silicon-based instead of carbon-based. This is unlikely for a number of reasons. Carbon when combined with oxygen forms the gas CO2. Silicon can too combine with oxygen to form SiO2, but the result is a solid and doesn’t dissolve in water, the chief solvent of life. Silicon also can’t form the elaborate chemicals of living things that carbon can build. Silicon dust, when inhaled, can cause lung disease.

What silicon is well suited for are computers, though we started out using something else in electronics. When John Bardeen, Walter Brattain, and William Shockley invented the transistor in the late 1940s, and won the Nobel Prize for their efforts, they employed germanium. But germanium is too temperamental an element, generating too much heat. By the mid-1950s the fledgling semiconductor industry switched to silicon.

Computers require lots and lots of transistors, but they are hard to solder onto a circuit board in such great numbers. In 1958 Jack Kilby of Texas Instruments developed the integrated circuit, which can be machined microscopically from a single block of material, the process automated and speedy. Kilby won a Nobel for his invention. 

During the mid-1800s in Heidelberg, Robert Bunsen studied arsenic, geysers, and volcanoes. He also developed the “Bunsen burner” (47) and invented the spectroscope, then heated up chemicals with the burner and sent the light they emit through the spectroscope’s prism, which split the light into unique spectral lines. With these new tools, scientists quickly began to identify new elements, and before long Bunsen’s students Dmitri Mendeleev and Julius Lothar Meyer constructed primitive periodic tables. The two won the prestigious Davy Prize.

Mendeleev’s chief insight was that elements can be defined by their atomic weight. He also predicted correctly that gaps in his chart would be filled by elements yet to be discovered and guessed what their weights and densities will be. Kean notes, “Overall, Mendeleev’s work is comparable to that of Darwin in evolution and Einstein in relativity. None of those men did all the work, but they did the most work, and they did it more elegantly than others” (53).

The chapter then turns to Paul-Émile Lecoq de Boisbaudran, who discovered in 1875 one of the missing elements to the chart. He named it gallium “after Gallia, the Latin name for France” (54), his home country. Mendeleev, however, wanted some of the credit. He alleged that Boisbaudran’s data was incorrect; when Boisbaudran redid his experiments, he found that Mendeleev was right. Thus, both men have a claim on the element’s discovery. Kean reveals the inspiration for his book’s title with an anecdote about gallium, which melts at 84 degrees Fahrenheit. A jokester can “fashion gallium spoons, serve them with tea, and watch as [the] guests recoil when their Earl Grey ‘eats’ their utensils” (54).

Kean now jumps back several centuries, discussing Marco Polo’s adventures to China. Polo brought Chinese porcelain back to Europe in the 1200s. Europeans tried to reproduce the bedeviling material, but their efforts proved fruitless. Two Polish craftsmen of the early 1700s finally discovered the secret: kaolin and feldspar cooked together at high temperatures.

Finally, Kean looks at a feldspar mine that opened in Ytterby, Sweden, in 1780. The quarry revealed unusual minerals alongside the expected ones. Geochemist Johan Gadolin analyzed the odd rocks and began to tease out the first of the lanthanides, or rare earths, that fit gaps in Mendeleev’s chart. Several were named for the place where they are first found: “ytterbium, yttrium, terbium, and erbium […] holmium, after Stockholm; thulium, after the mythic name for Scandinavia; and, at Lecoq de Boisbaudran’s insistence, Gadolin’s namesake, gadolinium” (62).