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The Disappearing Spoon
And Other True Tales of Rivalry, Adventure, and the History of the World from the Periodic Table of the Elements (Young Readers Edition)
By Sam Kean
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Why did Gandhi hate iodine (I, 53)? How did radium (Ra, 88) nearly ruin Marie Curie’s reputation? And why did tellurium (Te, 52) lead to the most bizarre gold rush in history?
The periodic table is a crowning scientific achievement, but it’s also a treasure trove of adventure, greed, betrayal, and obsession. The fascinating tales in The Disappearing Spoon follow elements on the table as they play out their parts in human history, finance, mythology, conflict, the arts, medicine, and the lives of the (frequently) mad scientists who discovered them.
Adapted for a middle grade audience, the young readers edition of The Disappearing Spoon offers the material in a simple, easy-to-follow format, with approximately 20 line drawings and sidebars throughout. Students, teachers, and burgeoning science buffs will love learning about the history behind the chemistry.
GEOGRAPHY OF THE ELEMENTS
WHEN YOU THINK OF THE PERIODIC TABLE, YOU PROBABLY THINK OF THAT COLORFUL chart with many columns and rows hanging on the wall of your science classroom. You may have talked about it in class, and you may even have been able to use it during tests and exams. Unfortunately, even when you could use it, this gigantic cheat sheet may have seemed less than helpful! But the table and each box on it are full of secrets waiting to be decoded.
On the one hand, the periodic table seems beautifully organized, but on the other, it can sometimes look like a jumble of long numbers, abbreviations, and what appear to be computer error messages ([Xe]6s24f15d1). So what does it look like with all the clutter removed? Sort of like a castle, with an uneven main wall and tall turrets on both ends. It has eighteen vertical columns and seven horizontal rows, with two extra rows below.
The castle is made of “bricks,” but the bricks are not interchangeable. Each brick shows an element, or type of substance (as of now, it has 118 officially named elements, with a few more coming soon), and together they make up the table. If any of those bricks didn’t sit exactly where it was supposed to, the entire castle would crumble. That’s no exaggeration: If scientists determined that one element somehow fit into a different slot, or that two of the elements could be swapped, the entire “castle” would tumble down. They all fit together in a particular way.
Seventy-five percent of the bricks are metals, which means most elements are cold gray solids, at room temperature. A few columns on the right contain gases. Only two elements, mercury (element 80) and bromine (element 35), are liquids at room temperature. In between the metals and gases, about where Kentucky sits on a map of the United States, lie some wacky elements that have crazy properties, such as the ability to make acids billions of times stronger than anything locked up in your school’s chemical supply room.
WHAT EXACTLY IS AN ELEMENT?
The term “elements” goes back to ancient Greece. There, the philosopher Plato came up with the word (in Greek, stoicheia). Of course, Plato didn’t know what an element really was in chemistry terms—he was using it to refer to air, water, earth, and fire.
Helium (element 2) is a good example of “element-ness”—a substance that cannot be broken down or altered by normal, chemical reactions. Today we say that carbon dioxide, for instance, isn’t an element because one molecule of carbon dioxide divides into carbon (element 6) and oxygen (element 8). But carbon and oxygen are elements because you cannot divide them without destroying them.
It took scientists twenty-two hundred years to finally work out what elements actually are, simply because it was hard to see what made carbon carbon when it appeared in thousands of different compounds, all with different properties. It’s kind of like the difference between chocolate ice cream and a chocolate chip cookie. They are both made with chocolate, but they are different in every other way (although they are both usually delicious). Nearly all elements form bonds with other elements to make compounds, and that makes it difficult to see what the pure elements themselves are really like. Scientists might have figured out what elements are much sooner had they known about helium, which exists only as a pure element and not in any compounds.
Helium acts this way for a reason. Each element is made up of a specific type of atom. All atoms contain negative particles called electrons, which live in different energy levels inside the atom. Each level needs a certain number of electrons to fill itself and become complete. In the innermost level, that number is two. Other numbers are required in other levels, but it’s often eight. Elements have equal numbers of negative electrons and positive particles called protons, so they’re electrically neutral because the positive and negative charges cancel out. Electrons, however, can be swapped between atoms, and when atoms lose or gain electrons, they form charged particles called ions.
IN 1911, A DUTCH-GERMAN scientist was cooling mercury with liquid helium when he discovered that below −452°F the system lost all electrical resistance and became an ideal conductor. This would be sort of like cooling an iPhone down to hundreds of degrees below zero and finding that the battery remained fully charged until infinity, as long as the helium kept it cold. A Russian-Canadian team pulled an even neater trick in 1937 with pure helium. When cooled down to −456°F, helium turned into a superfluid, with exactly zero resistance to flow. Superfluid helium defies gravity and flows uphill and over walls. Not even Plato would have predicted something so “cool” could actually happen in real life!
Electrons are arguably the most important part of an atom. They take up virtually all of an atom’s space, like clouds swirling around the nucleus, an atom’s tiny core. If an atom were blown up to the size of a football stadium, the nucleus would be a tennis ball at the fifty-yard line.
What’s important to know is that atoms fill their inner, lower-energy levels as full as possible with their own electrons, but when they undergo chemical reactions, they either lose, gain, or share electrons to secure the right number in the outermost level. Helium has exactly the number of electrons it needs to fill its only level, so there is no need for it to interact with other atoms or to lose, gain, or share electrons. This makes helium tremendously independent, possibly even “noble.”
Electrons drive the periodic table, and no one did more to explain how electrons behave than American chemist Gilbert N. Lewis. Lewis spent his whole life working out how an atom’s electrons work, especially in acids, and in their chemical opposites, bases. But one of the things he’s best known for is that he was probably one of the greatest scientists never to win the Nobel Prize—and he was pretty bitter about it. Part of the reason he never won it is that he didn’t discover anything that you could point to and say, “Wow! Look at that amazing thing!” Instead, he spent his life refining our understanding of how electrons work in many contexts, something that would help future scientists in a big way.
Before about 1890, scientists judged acids and bases by tasting or dunking their fingers into them—er… not a great idea or a very scientific one, either! Pretty soon, scientists realized that many acids contain hydrogen (element 1), the simplest element, which consists of just one electron and one proton. When an acid like hydrochloric acid (HCl) mixes with water, it splits into H+ and Cl− ions. Removing the negative electron from the hydrogen atom leaves just a bare proton, the H+. Weak acids like vinegar pop a few H+’s into the solution, while strong acids like sulfuric acid flood solutions with them.
Lewis decided this definition of an acid limited scientists too much, since some substances act like acids without relying on hydrogen. Instead of saying that H+ splits off, he emphasized that Cl− runs away with hydrogen’s electron, like an electron thief. In contrast, bases (which, remember, are the chemical opposites of acids) may be called electron givers or donors. These so-called Lewis definitions of acids as electron pair acceptors and bases as electron pair donors emphasize the importance of electrons, and it fits better with the electron-dependent chemistry of the periodic table.
Although Lewis’s acid theory is almost one hundred years old, scientists are still using his ideas to make stronger and stronger acids. You may know that acid strength is measured by the pH scale, with lower numbers on the scale like 1, 2, and 3 meaning stronger acids, and higher numbers like 12, 13, and 14 meaning stronger bases. In 2005, a chemist from New Zealand invented a boron (element 5)–based acid called a carborane, with a pH of –18 (yes, negative eighteen!). To put that in perspective, water has a pH of 7, and the concentrated HCl in our stomachs has a pH of 1. Because of the weird math of the pH scale, dropping one unit (e.g., from pH 4 to pH 3) boosts an acid’s strength by ten times. So moving from stomach acid, at pH 1, to the boron-based acid, at pH –18, means that the boron-based acid is ten billion billion times stronger than HCl.
There are even stronger acids based on antimony (element 51), which has a colorful history. Nebuchadnezzar, a Babylonian king in the sixth century bc, unknowingly used a poisonous antimony-lead mix to paint his palace walls yellow. Perhaps not coincidentally, he soon went mad, sleeping outdoors in fields and eating grass like an ox. Later, around the nineteenth century, antimony pills were used as laxatives. Unlike modern pills, these hard pills didn’t dissolve in the stomach, and they were considered so valuable that people rooted through poop to find and reuse them. Yuck. Some lucky families even passed (!) pills from father to son!
In fact, antimony was often used in medicine, since people didn’t yet realize how toxic it was. Mozart probably died from taking too much of it to combat a severe fever.
We’ve come a long way in our understanding of acids and bases and how electron behavior drives the periodic table. But to really understand the elements, you can’t ignore the part that makes up more than 99 percent of their mass—the nucleus. Whenever atoms react chemically, the nucleus remains unchanged, and only the electrons matter. Inside the nucleus, the number of positive protons—the atomic number, the whole number that’s usually somewhere above the letters in each box in the periodic table—determines the atom’s identity. In other words, an atom of one element cannot gain or lose protons without becoming an atom of an entirely different element.
And whereas electrons obey the laws of the greatest scientist never to win the Nobel Prize, the nucleus obeys the rules of probably the most unlikely Nobel laureate ever, Maria Goeppert-Mayer.
Maria Goeppert was born in Germany in 1906. Even though her father was a sixth-generation professor, Maria had trouble persuading a PhD program to admit a woman, so she bounced from school to school, taking lectures wherever she could. When she finally earned her doctorate, no university would hire her. She could enter science only through her husband, Joseph Mayer, an American chemistry professor visiting Germany. She returned to Baltimore with him in 1930, and the newly named Goeppert-Mayer began tagging along with Mayer to work and conferences. Unfortunately, Mayer lost his job several times during the Great Depression, and the family drifted to universities in New York and then Chicago.
Most schools tolerated Goeppert-Mayer hanging around to chat science. Some even condescended to give her work, though they refused to pay her, and the topics were stereotypically “feminine,” such as figuring out what causes colors. After World War II, the University of Chicago finally took her seriously enough to make her a professor of physics. Although she got her own office, the department still didn’t pay her.
Eventually she and her husband moved to a new university in San Diego that actually paid her a salary. By then she had discovered something called the nuclear shell model, which helped scientists understand the structure of the nucleus—but she still hadn’t discovered a way to make everyone take her seriously as a scientist. When the Swedish Academy announced in 1963 that she had won her profession’s highest honor, her local San Diego newspaper greeted her big day with the headline “S.D. Mother Wins Nobel Prize.”
LOCATION, LOCATION, LOCATION
The position of each element on the table, which is determined by its atomic number (i.e., the number of protons), is crucial—its geography determines nearly everything scientifically interesting about it. So, in addition to visualizing it as a castle, you may think of the periodic table as a map.
First up, in column eighteen, at the far right-hand side, is a set of elements known as the noble gases. Many chemists find noble gases fascinating and, as with the idea of elements, we can trace this fascination with noble gases back to Plato. For someone who knew nothing about chemistry, he definitely had a big impact on it. If Plato had known what elements actually were, he might have selected the elements on the eastern edge of the table, especially helium, as his favorites.
Why? Well, in his writings, Plato said that unchanging things are more “noble” than things that interact with others. Helium and the rest of the noble gases tend not to react with other things, so Plato would probably have loved them.
Helium isn’t the only element that has exactly the number of electrons it needs. The same idea extends down the entire eighteenth column beneath it—the gases neon (element 10), argon (element 18), krypton (element 36), xenon (element 54), and radon (element 86) all have the electrons that they need, so none of them reacts with anything under normal conditions.
The behavior of the noble gases is rare, however. One column to the west sits some of the most energetic and reactive elements on the periodic table, the halogens. Even more violent elements appear on the western edge, the alkali metals.
In addition to the reactive alkali metals on its west coast and halogens and noble gases up and down its east coast, the periodic table contains a “great plains” that stretches right across its middle—columns three through twelve, the transition metals.
As we move horizontally across the periodic table, each element has one more electron than its neighbor to the left. Sodium (element 11) normally has eleven electrons; magnesium (element 12) has twelve electrons; and so on. The addition of one electron to each transition metal would normally alter its behavior, as happens with elements in other parts of the table. But not those pesky transition metals. Chemically, many transition metals look and behave similarly. That’s because, instead of exposing their outer electrons to the world (the way most elements do), transition metals often hide their outer electrons in a sort of secret compartment. As a result, transition metals tend to look the same to the outside world and behave the same way in chemical reactions.
Despite being normal metals in some ways, the alkali metals, instead of slowly rusting or corroding, can spontaneously combust in air or water. They also react easily with the halogens. The halogens have seven electrons in their outer layer, one short of the eight that they need, while the alkalis have one electron in their outer level and a full set in the level below. So it’s natural for the group one metals to dump their extra electron on the group seventeen halogens and for the resulting positive and negative ions to form strong links. (When it comes to ions, opposites attract: Positive and negative ions are drawn to each other like magnets.)
This sort of linking—called ionic bonding—explains why combinations of halogens and alkali metals, such as sodium chloride (table salt, NaCl), are common. It’s the easiest way for all atoms to get the electrons that they need. In a similar way, two ions of sodium (Na+) take on one of oxygen (O2-) to form sodium oxide (Na2O). Overall, you can usually tell at a glance how elements will combine by noting their column numbers and figuring out their charges. Unfortunately, not all of the periodic table is so clean and neat. But the weird behavior of some elements actually makes them even more interesting.
THE FATHERS OF THE PERIODIC TABLE
YOU MAY SAY THE HISTORY OF THE PERIODIC TABLE IS REALLY THE HISTORY OF THE many people who shaped it. Just as some of the elements of the periodic table are better known than others, many of the names of the scientists who discovered them and who arranged them into the first periodic tables are famous, while others have long since been forgotten.
BUNSEN, MENDELEEV, AND MEYER
One name from the periodic table’s history that you may recognize is Robert Bunsen. This pioneer of the periodic table deserves special praise, since you’ve probably used a piece of school lab equipment named after him. Disappointingly, the German chemist didn’t actually invent “his” Bunsen burner, but rather improved on the original design in the mid-1800s.
Bunsen’s first love was arsenic, one of the most poisonous elements on the periodic table. Although element 33 has had quite a reputation since ancient times (Roman assassins used to smear it on figs and wait for their targets to eat), few law-abiding chemists knew much about arsenic before Bunsen started sloshing it around in test tubes. He worked mainly with arsenic-based cacodyls, chemicals whose name is based on the Greek word for “stinky.” Cacodyls smelled so foul, Bunsen said, that they made him hallucinate. His tongue became “covered with a black coating.” Perhaps from self-interest, he soon developed what is still the best antidote to arsenic poisoning, iron oxide hydrate, a chemical related to rust that clamps onto arsenic in the blood and drags it out. Still, he couldn’t shield himself from every danger. The accidental explosion of a glass beaker of arsenic nearly blew out his right eye, leaving him half-blind for the last sixty years of his life.
Following the accident, Bunsen put arsenic aside and, after investigating geysers and volcanoes for a while, he settled back into chemistry at the University of Heidelberg in the 1850s. There, he invented the spectroscope, a piece of lab equipment that uses a prism and light to study elements. Each element on the periodic table produces sharp, narrow bands of colored light when heated. Hydrogen, for example, always emits one red, one yellowish-green, one baby-blue, and one indigo band. If you heat some mystery substance and it emits those specific lines, you can bet it contains hydrogen. This was a powerful breakthrough, as it was the first way to peer inside exotic compounds without boiling them down or disintegrating them with acid.
The only thing limiting spectroscopy at that point was getting flames hot enough to excite elements. So Bunsen created the device that made him a hero to everyone who ever melted a ruler or set a pencil on fire in a school chemistry lab!
Bunsen’s work helped the periodic table develop rapidly for two reasons: First, the spectroscope allowed new elements to be identified. Second, and just as important, it helped sort through many claims for new elements by finding old elements in disguise in these unknown substances.
But beyond finding new elements, scientists needed to organize them into a family tree of some sort. And here we come to Bunsen’s other great contribution to the table. At Heidelberg, he instructed a number of people responsible for early work in periodic law. This includes our second character, Dmitri Mendeleev, the man often referred to as the father of the periodic table. All the scientists working on early periodic tables recognized likenesses among certain elements. But some scientists were better than others at recognizing those similarities. Knowing how to recognize and predict such similarities soon enabled Mendeleev to create the first real periodic table.
- On Sale
- Sep 10, 2019
- Page Count
- 240 pages
- Little, Brown Books for Young Readers