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A Visual Exploration of Every Known Atom in the Universe
By Nick Mann
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This item is a preorder. Your payment method will be charged immediately, and the product is expected to ship on or around April 3, 2012. This date is subject to change due to shipping delays beyond our control.
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Includes a poster of Theodore Gray’s iconic photographic periodic table of the elements!
Based on seven years of research and photography by Theodore Gray and Nick Mann, The Elements presents the most complete and visually arresting representation available to the naked eye of every atom in the universe. Organized sequentially by atomic number, every element is represented by a big beautiful photograph that most closely represents it in its purest form. Several additional photographs show each element in slightly altered forms or as used in various practical ways. Also included are fascinating stories of the elements, as well as data on the properties of each, including atomic number, atomic symbol, atomic weight, density, atomic radius, as well as scales for electron filling order, state of matter, and an atomic emission spectrum.
This of solid science and stunning artistic photographs is the perfect gift book for every sentient creature in the universe.
How the Periodic Table Got Its Shape
Hang on tight, we’re going to explain quantum mechanics in one page. (If you find this section too technical, feel free to skim it—there isn’t going to be a quiz at the end.)
Every element is defined by its atomic number, the number of positively charged protons in the nucleus of every atom of that element. These protons are matched by an equal number of negatively charged electrons, found in “orbits” around the nucleus. I say “orbits” in quotes because the electrons are not actually moving around their orbits like planets around a star. In fact, you can’t really speak of them as moving at all.
Instead, each electron exists as a probability cloud, more likely to be in one place than another, but not actually in any one place at any given time. The figures below show the various three-dimensional shapes of the probability clouds of electrons around a nucleus.
The first type, called an “s” orbital, is totally symmetrical—the electron is not any more likely to be in one direction than another. The second type, called a “p” orbital, has two lobes, meaning the electron is more likely to be found on one side or the other of the nucleus, and less likely to be found in any direction in between.
While there is only one “s”-type orbital, there are three “p” types, with lobes pointing in the three orthogonal directions (x, y, z) of space. Similarly there are five different types of “d” orbitals and seven different types of “f” orbitals, with increasing numbers of lobes. (You may think of these shapes as a bit like three-dimensional standing waves.)
Each shape of orbital can exist in multiple sizes, for example the 1s orbital is a small sphere, 2s is a larger sphere, 3s is larger still, and so forth. The energy required for an electron to be in any given orbital increases as the orbit becomes bigger. And all else being equal, electrons will always settle into the smallest, lowest-energy orbit.
So do all the electrons in an atom normally sit together in the lowest-energy 1s orbital? No, and here we come to one of the most fundamental discoveries in the early history of quantum mechanics: No two particles can ever exist in exactly the same quantum state. Because electrons have an internal state known as “spin,” which can be either up or down, it turns out that exactly two electrons can reside in a given orbital—one with spin up and one with spin down.
Hydrogen has only one electron, so it sits in the 1s orbital. Helium has two, and they both fit into 1s, filling it to its capacity of two. Lithium has three, and since there is no room in 1s anymore, the third electron is forced to sit in the higher-energy 2s orbital. And so on—the orbitals are filled one at a time in order of increasing energy.
Look at the Electron Filling Order diagram on the right side of any element page in this book, and you’ll see a graph of the possible orbitals from 1s to 7p, with a red bar indicating which ones are filled with electrons (7p is the orbital of highest energy occupied by electrons of any known element). The exact order in which orbitals are filled turns out to be surprisingly subtle and complex, but you can watch it happen as you flip through the pages of this book. Pay particular attention around gadolinium (64)—if you think you’ve got it figured out, your confidence might be shaken by what happens there.
It is this filling order that determines the shape of the periodic table. The first two columns represent electrons filling “s” orbitals. The next ten columns are electrons filling the five “d” orbitals. The final six columns are electrons filling the three “p” orbitals. And last but not least, the fourteen rare earths are electrons filling the seven “f” orbitals. (If you’re asking yourself why helium, element 2, is not above beryllium, element 4, congratulations—you’re thinking like a chemist rather than a physicist. Eric Scerri’s book, referenced in the bibliography, is a good start toward answering such questions.)
Everything you need to know. Nothing you don’t.
The mini table on every element page has one highlighted yellow square to show you where that element is located on the periodic table. The colors divide the table into the groups described on the preceding pages.
An element’s atomic weight (not to be confused with its atomic number) is the average weight per atom in a typical sample of the element, expressed in “atomic mass units,” or amu. The amu is defined as 1/12 the mass of a 12C atom. Roughly speaking, one amu is the mass of one proton or one neutron, and thus an element’s atomic weight is approximately equal to the total number of protons and neutrons in its nucleus.
However, you will notice that the atomic weights of some elements fall well between whole integers. When typical samples of an element contain two or more naturally occurring isotopes, the averaging of isotopic weights explains the fractional amu. (Isotopes are explained in more detail under protactinium, element 91; the basic idea is that an element’s isotopes all have the same number of protons, and thus the same chemistry, but differ in the numbers of neutrons in their nuclei).
The density of an element is defined as the idealized density of a hypothetical flawless single crystal of the absolutely pure element. This can never be realized exactly in practice, so the densities are generally calculated from a combination of the atomic weight and x-ray crystallographic measurements of the spacing of atoms in crystals. The density is given in units of grams per cubic centimeter.
The density of a material depends on two things: how much each atom weighs, and how much space each atom takes up. The atomic radius shown for each element is the calculated average distance to the outermost electrons from the nucleus in picometers (trillionths of a meter). The diagrams are merely schematic—they represent all the electrons in their respective electron shells, with the overall size matching the size of the atom, but the position of individual electrons is not to scale, nor do electrons actually exist as sharp points spinning around the atom. The dashed blue reference circle shows the radius of the largest of all atoms, cesium (55).
The crystal structure diagram shows the arrangement of atoms (the unit cell that is repeated to form the whole crystal) when the element is in its most common pure crystalline form. For elements that are normally gas or liquid, this is the crystal form they take on when they are cooled enough to freeze solid.
ELECTRON FILLING ORDER
This diagram shows the order in which electrons fill the available atomic orbitals, which are explained in detail on the preceding page.
ATOMIC EMISSION SPECTRUM
When atoms of a given element are heated to very high temperatures, they emit light of characteristic wavelengths, or colors, which correspond to the differences in energy levels between their electron orbitals. This diagram shows the colors of these lines, each one corresponding to a particular energy-level difference, arranged into a spectrum from the barely visible red at the top to the nearly ultraviolet at the bottom.
STATE OF MATTER
This temperature scale in degrees Celsius shows the range of temperatures over which the element is solid, liquid, or gas. The boundary between solid and liquid is the melting point, while the boundary between liquid and gas is the boiling point. Twist the pages of the book to spread the edges of the pages out, and you will see a graph of the melting and boiling points, which shows very pronounced trends across the periodic table.
Stars shine because they are transmuting vast amounts of hydrogen into helium. Our sun alone consumes six hundred million tons of hydrogen per second, converting it into five hundred and ninety-six million tons of helium. Think about it: Six hundred million tons per second. Even at night.
And where does the other four million tons per second go? It’s converted into energy according to Einstein’s famous formula, E=mc2. About three-and-a-half-pounds-per-second’s worth finds its way to the earth, where it forms the light of the dawn rising, the warmth of a summer afternoon, and the red glow of a dying day.
The sun’s ferocious consumption of hydrogen sustains us all, but hydrogen’s importance to life as we know it begins closer to home. Together with oxygen it forms the clouds, oceans, lakes, and rivers. Combined with carbon (6), nitrogen (7), and oxygen (8), it bonds together the blood and body of all living things.
Hydrogen is the lightest of all the gases—lighter even than helium—and much cheaper, which accounts for its ill-advised use in early airships such as the Hindenburg. You may have heard how well that went, though in fairness the people died because they fell, not because they were burned by the hydrogen, which in some ways is less dangerous to have in a vehicle than, say, gasoline.
Hydrogen is the most abundant element, the lightest, and the most beloved by physicists because, with only one proton and one electron, their lovely quantum mechanical formulas actually work exactly on it. Once you get to helium with two protons and two electrons, the physicists pretty much throw up their hands and let the chemists have it.
Helium is named for the Greek god of the sun, Helios, because the first hints of its existence were dark lines in the spectrum of sunlight that could not be explained by the presence of any elements known at the time.
It might seem a paradox that an element common enough to fill party balloons with was the first element to be discovered in space. The reason is that helium is one of the noble gases, so named because they do not interact with the common riffraff of elements, remaining inert and aloof to nearly all chemical bonding. Because it does not interact, helium could not easily be detected by conventional wet chemical methods.
As a replacement for hydrogen in airships, helium, which is completely nonflammable, has much to recommend it. The main problem is that it’s a lot more expensive, and provides somewhat less lift. Anyone want to go for a ride in the low bid model?
The helium we use today is extracted from natural gas as it comes out of the ground. But unlike all other stable elements, it was not deposited there when the earth was formed. Instead it was created over time by the radioactive decay of uranium (92) and thorium (90). These elements decay by alpha particle emission, and “alpha particle” is simply the physicist’s name for the nucleus of a helium atom. So when you fill a party balloon, you’re filling it with atoms that just a few tens or hundreds of millions of years ago were random protons and neutrons in the nuclei of large radioactive atoms. That, frankly, is weird. Though not as weird as the way lithium messes with your mind.
Lithium is a very soft, very light metal. So light that it floats on water, a feat matched by only one other metal, sodium (11). While floating on water, lithium will react with that water, releasing hydrogen gas at a steady, moderate rate. (The real excitement in this department begins with sodium.)
Despite its reactive nature, lithium is widely used in consumer products. Lithium metal inside lithium-ion batteries powers countless electronic devices, from pacemakers to cars, including the laptop on which I am typing this text. Lithium-ion batteries pack tremendous power into not much weight, in part because of lithium’s low density. Lithium stearate is also used in the popular lithium grease found on cars, trucks, and mechanics.
People who pay attention to these things have noticed an interesting fact: There’s only one place in the world with a really large amount of easily recoverable lithium. If electric cars based on lithium-ion batteries ever become very widespread, you might want to keep an eye on Bolivia.
The lithium ion has another trick up its sleeve: It keeps some people on an even emotional keel. For reasons that are only vaguely understood, a steady dose of lithium carbonate (which dissolves into lithium ions in the body) smoothes out the highs and lows of bipolar disorder. That a simple element could have such a subtle effect on the mind is testimony to how even a phenomenon as complex as human emotion is at the mercy of basic chemistry.
Lithium is soft, reactive, and helps keep things in balance. Beryllium is, well, let’s just say different.
Beryllium is a light metal (though three and a half times the density of lithium, it’s still significantly less dense than aluminum, element 13). Where lithium is soft, low-melting, and reactive, beryllium is strong, melts at a high temperature, and is notably resistant to corrosion.
These properties, combined with its high cost and poisonous nature, account for the unique niche beryllium has carved out for itself: missile and rocket parts, where cost is no object, where strength without weight is king, and where working with toxic materials is the least of your worries.
Beryllium has other fancy applications. It is transparent to x-rays, so it’s used in the windows of x-ray tubes, which need to be strong enough to hold a perfect vacuum, yet thin enough to let the delicate x-rays out. A few percent of it alloyed with copper (29) forms a high-strength, nonsparking alloy used for tools deployed around oil wells and flammable gases, where a spark from an iron tool could spell disaster, in great big flaming red letters.
In keeping with the sport of golf’s tendency to use high-tech materials out of a desperate hope that they may help get the ball where it’s supposed to go, beryllium copper is also used in golf-club heads. Needless to say, it doesn’t help any more than the manganese bronze or titanium (22) used for the same purpose.
Combining beauty with brawn, the mineral beryl is a crystalline form of beryllium aluminum cyclosilicate. You may be more familiar with the green and blue varieties of beryl, which are known as emerald and aquamarine.
Beryllium: A debonair, James Bond–style metal able to launch rockets one minute and charm the ladies the next. Then there’s boron.
Poor boron—with a name like that, how can it get any respect? It doesn’t help that boron’s most commonly found in borax, the laundry aid. But boron is more glamorous than you might think.
Combine boron (5) with nitrogen (7), and you get crystals similar to those of their average, carbon (6), the element that forms diamond. Cubic boron nitride crystals are very nearly as hard as diamond, but much less expensive to create and more heat resistant, making them popular abrasives for industrial steelworking.
Recent theoretical calculations indicate that the alternate wurtzite-crystal form of boron nitride, as yet never created in single-crystal form, might actually be harder than diamond under certain conditions, and for certain technical definitions of “hard.” Unseating diamond from its long reign as the hardest known material would be quite a coup, but for the time being “wurtzite” boron nitride’s only accomplishment is causing an annoying footnote you now have to put next to any claim that diamond is the hardest known substance.
Boron carbide, also one of the hardest known substances, even has a genuine secret-agent application: Granules of it poured into the oil-fill hole of an internal combustion engine will destroy the engine by irreparably scoring the cylinder walls. Of slightly less interest to the CIA is the fact that boron is critical in cross-linking the polymers that gives Silly Putty its amazing ability to be both soft and moldable in your hand, yet hard and bouncy when you throw it against the wall.
But while boron is not quite the frump you might expect from its name, it’s really not in the same league as carbon.
Carbon is the most important element of life, period. Sure, there are many others without which life would not exist, but from the spiral backbone of DNA to the intricate rings and streamers of the steroids and proteins, carbon is the element whose unique properties tie it all together. The very term “organic compound” refers exclusively to chemicals containing carbon.
Not content to be the foundation of all life on earth, carbon also forms diamond, the hardest known substance (at least for now; challengers are discussed under boron, element 5). But contrary to popular belief, diamonds are not particularly rare, nor are they unusually beautiful, nor are they forever: all three are myths created by the DeBeers diamond company. Diamonds would cost a tenth as much but for DeBeers’s monopoly control. Cubic zirconia or crystalline silicon carbide are just as pretty. And at high enough temperatures, diamonds burn up into nothing but carbon dioxide.
If I were writing these words twenty-five years or so ago, I would probably have been doing it with carbon. The “lead” in pencils is actually graphite, a form of carbon, and has been since the 16th-century discovery in the English Lake District of the great mine at Borrowdale, the first source of pure graphite.
Carbon atoms like to form sheets, like a honeycomb with a carbon atom at each corner. Stack the sheets and you have graphite. Fold them into a sphere and you have a C60 “buckyball,” named for Buckminster Fuller who invented the geodesic dome. Roll the sheets into tubes and you have the strongest material known to science: carbon nanotubes.
Carbon has now become a focus of political controversy centered on the fact that our civilization is pumping carbon dioxide back into the atmosphere at about 100,000 times the rate it was put away by the dinosaurs and their swamps. Interestingly, the situation with nitrogen is exactly reversed.
At the same time that modern civilization has been pumping carbon dioxide into the atmosphere, we’ve been pulling out nitrogen and eating it.
Nitrogen as N2 in the air is inert and largely useless, but when it’s converted to a more reactive form, such as ammonia (NH3), it becomes a vital fertilizer. Only some plants, beans for example, aided by microorganisms residing in their roots, are able to draw the nitrogen they need directly from the air. This is one reason that, before the advent of cheap nitrogen fertilizer, corn, which cannot “fix” nitrogen, was alternated in the fields with beans or alfalfa, which leave the soil with more nitrogen than it started with.
Just before World War I, Fritz Haber invented a practical process for converting nitrogen from the air into ammonia, one of the most important discoveries in human history. Ammonia fertilizer now feeds a third of the world (the rest being fed mainly by phosphate fertilizers). His work with chlorine (17) was less benevolent, as you can read about under that element.
And since plant growth absorbs carbon dioxide from the air, nitrogen fertilization even helps, at least a bit, with alleviating the effects of global warming.
Liquid nitrogen is a cheap and readily available cryogenic cooling liquid. With a boiling point of -196°C it is cold enough to freeze almost anything. It is used to preserve biological samples, to amuse children by freezing and shattering flowers, and occasionally to make ice cream in record time.
There’s a lot of nitrogen around: Over 78 percent of the atmosphere is nitrogen. What’s the other 22 percent? Most of it is the oxygen we need to breathe.
- "Gray's trademark dry wit and historical anecdotes bring even the most basic lumps to life."—Popular Science
- "The Elements is a loving reimagination of the classic table."—Wired
- "I don't know if this is the first coffee-table book paying lush photographic homage to the periodic table, but it is certainly the most gorgeous one I've seen."—John Tierney, The New York Times
- On Sale
- Apr 3, 2012
- Page Count
- 240 pages
- Black Dog & Leventhal