Elemental

by Tim James

A table is supposed to be a neat rectangle with columns and rows. You know, like the one Lavoisier created. The periodic table we use today looks like a chimp accidentally vacuumed up a computer keyboard and tried to glue it back together with silly putty. It doesn’t look table-ish at all. So, who came up with the design and why did everyone else say “Yup, looks good to me”?

  The man to thank is one Alfred Werner, the Swiss Nobel Prize–winning chemist who published a short article in 1905 with the catchy title “Contribution to the development of a periodic system.”1 It was here that the periodic table first took shape.

  Let’s consider the first ten elements. Actually, let’s not. Let’s ignore elements 1 and 2 and begin with element 3. (I’ll explain in a moment.)

  We could line the elements up in a nice long row and be done with it:

  But now we can do better thanks to the Schrödinger equation. The first two elements of this row put their electrons into spherical orbitals while the next six go into dumbbell-shaped orbitals. This means we can split the line like so:

  The next eight elements have the same orbital shapes. The atoms will be bigger but they will otherwise have very similar chemistry. To represent this, we use Newlands’s periodic idea and add a second row to our table, still dividing into two blocks:

  Each column of elements represents a particular orbital shape. The only difference is that, as we go down, the orbitals get larger.

  When we get to element 21 a new shape gets introduced (quantum mechanics is like that). The outer electrons of the atom of this element, scandium, up to element 30, zinc, are shaped like bundles of balloons rather than dumbbells so we need to introduce a new block to the table. Element 31 goes back to the dumbbell shape and so our table now looks like this:

  It’s a bit irritating that nature insists on introducing weird orbitals when we get to larger elements, but that’s why the table is an awkward shape. It’s because nature is.

  Now, if you read from left to right across a row (period) you’re reading in ascending proton number, while the column (group) tells you what shape the atoms are going to have. Reach the end of one period and you just go on down to the next one.

  When Alfred Werner included all the known elements and orbital shapes, the table ended up like this:

  Suppose you wanted to know about iodine. By counting from left to right you learn that it is element number 53, meaning it will have fifty-three protons and fifty-three electrons. You can see it’s in the right-hand block (dumbbell shaped) so you also know what angles it will make with other atoms.

  Directly above it are chlorine and fluorine, both colorful non-metals. Iodine is in the same column so it will probably be a colorful non-metal too, but with a higher density as it’s on a lower period. Sure enough, we discover that these are exactly the properties of iodine.

  You can even use the periodic table to predict the properties of elements nobody has ever seen. Directly below iodine is astatine, the rarest element in the Earth’s crust (less than 1 g exists on the whole planet), but if we had a sample it would probably behave like a denser version of iodine. God bless quantum mechanics.

  ARCHITECTURAL SIMPLICITY

  You probably know from your periodic table T-shirt, mouse pad, shower curtain, pencil case, and notebook (I own all of these and assume everyone else does too) that the above pictures aren’t quite there yet.

  This fully expanded version of the table is rather cumbersome, so for simplicity we take one of the blocks, scooch it down to the bottom, and slide the others in to meet each other.

  This form of the periodic table was proposed by Glenn Seaborg in

  1945 and soon became the standard thanks to its simplicity and the fact that Seaborg did a lot of work to popularize science.2 But obviously we’ve missed out elements 1 and 2.

  Hydrogen and helium are both spherical atoms meaning they belong in groups 1 and 2 respectively:

  That’s what Harvey White did with his periodic table design in 1934 and it’s what Schrödinger would have wanted.3 Unfortunately, due to their small size, H and He don’t behave quite the same as the other elements in that block.

  They actually have more in common with elements on the other side of the table so if we placed them according to reactivity we would end up with things looking like this:

  This is what Ernst Riesenfeld did with his periodic table in 1928 (and what Mendeleev would have wanted).4

  Glenn Seaborg couldn’t make his mind up about where to put these two fiddly elements, sometimes drawing them on one side and sometimes the other (and briefly putting hydrogen in two groups in 1945).5

  Eventually, the general agreement was to acknowledge both the electron-orbital work of Schrödinger and the chemical-properties work of Mendeleev.

  So we split them up and put them at each end of the table, one for each scientist. It’s not logical but it’s a nice tribute to the two men who built it. And voilà, we have ourselves a periodic table.

  CHAPTER SEVEN

  Things that Go Boom

  THE MOST EXPLOSIVE EXPLOSIVE

  With the exception of nuclear weapons, most explosives work in the same way. First, a material is synthesized, which is highly unstable. In chemical terms this means it will fall apart given the chance. Second, the material is provoked, giving it the chance to break up and rearrange into stable substances. During this rearrangement a whack-ton of energy gets released (whack-ton being the technical term) in the form of light and heat.

  In addition, a small amount of solid or liquid explosive will expand rapidly to become a large volume of gas. This sudden expansion combined with lots of heat and light is what we call an explosion.

  Some substances are so unstable that even a little bit of agitation will cause the reaction to begin. Gunpowder only needs a candle flame to decompose, while TNT requires nothing more than a spark. You can even buy bang snaps, children’s toys consisting of small paper parcels filled with silver fulminate, a chemical that explodes when struck. The parcel is thrown at the ground and the impact causes a loud snap.

  Fireworks work on a similar principle. A powdered metal is launched into the air and once the fuse within has burned, a detonator gel will trigger, converting it into a gas. Each fleck of powder is sprayed outward as the gas expands, becoming so hot they start reacting with the atmospheric oxygen, giving us sparks.

  We’ve already seen in Chapter 4 that different elements give out different types of light, so by picking particular metals we get particular spark colors. Sodium will turn yellow, barium goes green, copper blue, and strontium brick red. Purple fireworks are famously difficult to achieve and usually involve a mixture of copper and strontium together.

  All explosives rely on the chemicals within being unstable, and the most unstable chemical ever created is called azidoazide azide, synthesized in 2011 by Thomas Klapötke. Although why anyone would want to make this chemical is beyond me.

  It contains fourteen nitrogen atoms and two carbon atoms clumped into branches around a tight ring, all squished together without much room. The bonds between the atoms are so strained that they spring apart under any circumstance.

  When Klapötke tried to dissolve the chemical in water, it exploded. When he tried to move it across his lab, it exploded. When he breathed in its general direction, it exploded. It even exploded when infrared light (the kind emitted from a TV remote control) was shone at it.1

  The best explosives, of course, are those that detonate on cue. They have to be stable enough to be moved, but unstable enough to still detonate. Azidoazide azide would be a poor choice, for the same reasons chlorine trifluoride was a terrible choice of rocket fuel. You’re better off with good old dynamite, invented by none other than Alfred Nobel.

  THE MERCHANT OF DEATH IS DEAD

  When somebody dies, we tend to say nice things about them because it’s taboo to speak ill of the dead. Such was not the case on February 12, 1888, when Alfred Nobel’s obituary was published. A French newspaper allegedly ran the headline
“The Merchant of Death is Dead” and went on to say, “Dr. Alfred Nobel, who became rich by finding ways to kill more people faster than ever before, died yesterday.”2

  A lot of people were not happy with a newspaper remembering the great scientist in such a way. That included Alfred Nobel himself, who managed to read his obituary owing to the fact that he was not really dead. The story goes that Alfred’s brother Ludvig had died and the newspaper mistook the two brothers, prompting them to publish their powerful attack.

  Nobel was a very talented chemist who had invented dynamite twenty-one years earlier. Initially he intended it for use in mining, but it had obvious military applications. Apparently, realizing what his legacy was, Nobel decided to alter his will and left his considerable fortune (then over thirty-one million Swedish kronor, equivalent to nearly eighty million US dollars today) as prize money to people who did things “for the greatest benefit of mankind.” The prizes were to be awarded for achievements in the three sciences, literature, and the promotion of peace: the Nobel Prizes.3

  The newspaper that ran the obituary is reported to be the Idiotie Quotidienne and I have desperately tried to find documentation to confirm its existence, but sadly I can find none.4 The fact the newspaper’s title translates to the Daily Idiocy might be a clue that the story is a hoax and indeed some of Nobel’s biographers have dismissed it as a persistent rumor.5

  It may be an apocryphal morality tale or perhaps an embellishment of Nobel’s reaction to his brother’s death. Whether it’s true or not, the Nobel Prizes are still considered the most prestigious awards it is possible to receive in science. The prize money is substantial, numbering in the millions, and it comes from Nobel’s estate built entirely on dynamite. Not literally, obviously. That would be stupid.

  The way dynamite works is simple. You take a large amount of silicon-based rock powder and soak it in a chemical called nitroglycerine. Pack this into a tube and stick a fuse in the end. As the fuse burns, the heat is transferred to the nitroglycerine-soaked powder and—boom!

  Nitroglycerine is one of the unstable compounds I mentioned earlier. Composed of carbon, nitrogen, oxygen, and hydrogen atoms, it is a chemical that will self-react, i.e. one nitroglycerine particle will react with another to produce a bunch of gases, mostly carbon dioxide and water.

  These gases expand to over twelve hundred times their original volume and reach a temperature of 5,000°C. The reaction is also fast with the expansion and heating taking place in under a microsecond.

  All of this comes down to the question we’re going to answer in this chapter—why do chemical reactions happen at all? What do we mean when we say a chemical is unstable and how do atoms bond in the first place? To make sense of it all, we’ll need to dive a little deeper into the quantum ocean.

  GIVING CHEMISTRY A BAD NAME

  The word chemistry comes from alchemy, but a better name for the subject would be electronics, because chemical reactions are all about electrons. The nucleus of an atom is tiny compared to the overall radius so it’s the electrons on the outside that are interacting with everything.

  And electrons are always on the move. Were an electron to cease moving it would simultaneously cease to exist because movement, like the charge it possesses, is part of an electron’s identity. A stationary electron exists no more than a four-sided triangle.

  So, if we take movement as a given, there are only two things an atomic electron can really do. It can move outward away from the nucleus or it can move inward toward it. These two behaviors underpin almost every chemical reaction you’ll meet.

  Let’s revisit the concept of orbitals that we got from the Schrödinger equation. They’re the regions around a nucleus where electrons spend their time.

  Orbitals are the permitted electron territories, but electrons aren’t confined to live their entire lives in the same one. They can hop around. When an electron jumps from one orbital to another it’s called a “quantum leap” and it can happen between any two orbitals, even ones that are empty.

  Obviously, electrons prefer to occupy an orbital near the nucleus because it carries the opposite charge but they don’t always get their way. If the innermost orbitals are inhabited, other electrons have to make do with being further out.

  An atom is really a hustling, bustling place where the orbitals nearest the nucleus are considered prime real estate and every electron wants to move in. If one of the inner electrons happens to vacate their orbital for some reason, an outer electron will quantum leap to replace it.

  These quantum leaps don’t happen at random, though. The rule that ultimately accounts for chemistry is as follows: if an electron absorbs a beam of light it gets bumped to an outer orbital and if it emits a beam of light it drops to an inner one.

  Some types of light have more ability to promote electrons while others have less. A blue beam can promote an electron to a far-out orbital while red light might only nudge it up by one level. In the same way, electrons from far-out orbitals have the ability to release blue light when they drop, while electrons already near the nucleus might only release red.

  This is how the fireworks and spectroscopy we’ve already mentioned work. Every atom has a unique orbital arrangement so every atom emits or absorbs a unique light spectrum. When the electrons start jumping from orbital to orbital, the distance they jump determines what kind of light gets emitted or absorbed depending on which way they’re traveling.

  The question everyone usually asks at this point is why electrons absorb or emit light in the first place. I’m afraid the answer is because that’s just the way nature is. It’s just one of the fundamental laws that were established during the big-bang expansion. The same way balls roll downhill as they obey the laws of gravity, electrons release and absorb light as they obey the laws of quantum mechanics.

  ABILITY AND STABILITY

  We’ve talked about some beams of light having more ability to promote electrons than others, and in science we sometimes replace the word ability with the word energy. I’ve mostly avoided using it up until now because it’s a word fraught with difficulty and misconception.

  People talk about energy as if it were a thing being transferred from place to place, but it isn’t really. You can’t hold a lump of energy but a lump of matter can possess the ability to bash into things or the ability to explode, i.e. it can possess energy.

  In the context of quantum chemistry, energy means “how capable a beam of light is to push an electron into a higher orbital.” You’ll sometimes hear scientists say that electrons in outer orbitals have “absorbed energy” and that this energy gets released when they drop. This is convenient shorthand but we have to be clear: it is light that gets absorbed and emitted. Light has the ability to promote electrons and therefore possesses energy, but energy is not an actual thing.

  The opposite of ability is what we mean by “stability” and it’s a measure of how much energy an electron has lost when it drops down, or how reluctant it is to shift upward from its present orbital.

  An electron from an inner orbital, close to the nucleus, is less willing to change because it is happy where it is. We describe it as being chemically “stable.” An electron in a higher orbital with a lot of energy (ability to release light) is very unstable, however, because it is not happy and will change given the chance.

  The diagram below shows what happens when an electron absorbs a beam of light. It jumps from a low-energy orbital to a high-energy orbital, becoming unstable.

  The next diagram shows the reverse process. This is a high-energy electron dropping down to a more stable orbital. The only difference is that light is being emitted here rather than absorbed.

  Ability and stability are always at odds with each other and govern an electron’s reactive behavior. Gaining energy means losing stability and vice versa. This trade-off between ability and stability is what determines whether a reaction will happen or not.

  SHAKE, RATTLE, AND ROLL

  Different kinds of light
will produce different kinds of effect on an atom. Infrared light, which is too low in energy to interact with the electrons in our eyes so we can’t see it, will cause the orbitals themselves to stretch and twist rather than shunting electrons between them. Microwaves do something similar except they cause the atom to spin, rather than twist and bend.

  If you beam atoms with infrared or microwave light the result is that the atoms start dancing around and bashing into each other, exchanging energy. Ultimately this happens via the same light-transfer mechanism (electrons on one atom release light to electrons on the other, promoting them to a higher orbital/making the atom twist or spin more) but it’s quicker and more convenient to talk about atoms colliding and transferring energy.

  These twists and spins of the atoms are what we call heat and it’s why you feel warm when infrared or microwave light hits your skin. It also provides the basis of microwave ovens by causing the water inside a piece of food to jiggle.

  Obviously, the hotter a sample of chemical, the more likely it is the atoms will bump into each other and trigger orbital rearrangements/twists/spins. Or, put another way, heating most reactions tends to make them happen faster.

  UNITED WE FALL

  Imagine being an electron tethered to an atom’s nucleus. If another atom approaches, its nucleus can draw you toward it at the same time. If the pull is strong enough you can be dragged into a position halfway between both nuclei and you are no longer occupying an atomic orbital but a “molecular orbital.” A molecular orbital is known by a more common name: a chemical bond.

  If the molecular orbitals are at lower energy than the atomic orbitals with which we started, then electrons on two approaching atoms can drop into a molecular orbital together, releasing light as they go. A bond between the atoms is formed and we have carried out a chemical reaction.