This article was published in Scientific American’s former blog network and reflects the views of the author, not necessarily those of Scientific American
The chemical elements—the varieties of atoms existing in nature and even some that are manmade—are an endless source of fascination. But something about them remains mysterious to most people, perhaps even to a lot of chemists.
I had a chance to remind myself of that conundrum while I reading two terrific books on the subject--The Disappearing Spoon, by Sam Kean, and Periodic Tales, by Hugh Aldersey-Williams--I was doing research to write up the text for ScientificAmerican.com’s Interactive Periodic Table.
The periodic table neatly arranges the elements into columns. Those elements that share a column are supposed to possess (roughly) similar chemical properties. Most people who know the table well and use it every day—which includes virtually every chemist, biochemist, materials scientist, and lots of engineers—hardly give it a thought, but the predictions of the table are some of the oddest facts in nature. The laws that govern most of the phenomena before our very eyes—everything from cellular metabolism to a laptop’s batteries—are dictated by the chemistry of the elements, but we rarely realize how weird and wonderful the underlying forces are.
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The significance of the table’s columns is most explicit in the elements on the leftmost column, called the alkali metals (which include lithium, sodium and potassium), and in those on the second column from the right, called the halogens (fluorine, chlorine and their pals). These two families comprise the most reactive of all elements. Atoms of halogens such as fluorine have an extreme tendency to rip an electron off of other atoms, while alkali atoms have an equally strong tendency to give an electron away. Thus, mix a halogen and an alkali metal and you get a perfect one-to-one match—a salt such as potassium chloride or sodium chloride, aka table salt.
More generally, the tendency is that of a progressive transition from electron-donation to electron-acceptance as you move from the left to the right on the periodic table. This gradient is quantitatively measured by the property called electronegativity, which (by and large) increases going from left to right. Once you reach the last column on the right, though, you reset the dial. The elements of that family are the noble gases which are virtually unreactive.
Every scientist will tell you that the reason why the families of chemicals that make up each column have more or less predictable behavior—in other words, the reason why there is a periodic table at all—is that all elements secretly “want to be noble gases.” In an atom of a noble gas such as neon (element 10), the orbits of the electrons around the nucleus are neatly and symmetrically arranged like the petals of a flower. The halogen fluorine (9), on the other hand, which has one electron fewer than neon, has a skewed orbital arrangement. It desperately “wants” that extra electron to make the arrangement symmetrical and happy. Meanwile, for the alkali metal sodium (11), the easiest thing to do to become like neon is to shed one electron.
It goes without saying that this narrative about atoms’ unfulfilled electronic desires is not the actual scientific explanation. The typical user of the periodic table doesn’t need to go much deeper: the true reasons have been figured out long time ago and the answer now lies buried in some difficult quantum physics textbook. But if we stop and think for a minute, we realize what an astoundingly bizarre phenomenon this is.
Take oxygen. The corroding—rusting, tarnishing, well, oxidating—chemical par excellence, oxygen reacts because it “wants to” complete its bouquet of outer electronic orbitals to a full set of eight, and thus mimic neon’s serene ikebana. It hardly matters if some of those orbitals also have to fly around other atoms: by sharing some of its electrons with two hydrogen atoms, for example, oxygen can form a molecule of water. As long as the outer orbitals form a full octet around its nucleus, chemical nirvana is within reach.
But here’s why this is so weird. If you take an oxygen atom in isolation and feed it an electron, it will capture it and become an oxygen ion O-. Now imagine that an extra electron flies by. The two objects, the electron and the oxygen ion, are both negatively charged, so they should repel each other. And yet, the oxygen craves that second electron so much that it will grab it and become a doubly ionized ion, O--. The reactivity of oxygen is so strong that it will win over the electrostatic (or Coulomb) repulsion, a force of notorious intensity. (The setting I am describing is, I admit, pretty abstract, as electrons and oxygen don’t really get a chance to interact in the vacuum, and probably the resulting ion would not be stable, but please bear with me.)
We often hear that everything that happens in nature—with the possible exception of dark energy, whose origin remains a complete mystery—is a result of one of four fundamental forces: gravity, the weak and strong nuclear forces, and electromagnetism. Chemical reactions and the formation of states of matter such as quasicrystals or people are entirely governed by the electromagnetic force. But electromagnetic interactions are varied and complex, and the electrostatic force is but one special case of them. Chemical reactions are electromagnetic, but most of the time they take place against, not following, electrostatic forces.
The mechanisms that power a battery, for example, have everything to do with the electromagnetic interactions we call chemistry; but contrary to what we often think, electrons do not move from the negative pole to the positive pole because of electrostatics. It’s not that the negative pole has a surplus of electrons—a negative charge—and the positive pole a depletion of them. Both electrodes in a battery are electrically neutral.
What really powers a battery is the difference in electronegativity between the materials its electrodes are made of. Take the voltaic pile, for example, the first battery in history, invented around 1800 by Alessandro Volta. The pile’s negative electrode is made of zinc (30) and the positive electrode is made of copper (29). Copper is slightly more electronegative than zinc*.
Thus, if you put the two metals next to each other (or if you connect them by a wire), some electrons will move from the zinc to the copper.
We say that copper is the positive pole and zinc is the negative one, but in reality, the transition of electrons will happen against electrostatic forces, not following them: the positive electrode, copper, will become negatively charged from the extra electrons, at the expense of the negative electrode, zinc which will charge positively!
Given such state of affairs, just connecting two materials made of atoms with different electronegativity (in the case of materials, the corresponding term is electrochemical potential) would not give out much in terms of energy. The negative charges accumulating on the positive pole would quickly become too strong and they would repel further electrons; the negative electrode, meanwhile, would quickly become positively charged from the loss of electrons and thus it would hang on more strongly to its remaining electrons by electrostatic attraction. The transfer of electrons would quickly come to a halt.
That’s why a battery is not made just of two electrodes, but includes an electrolyte. The electrolyte permits the transfer of ions from the negative pole to the positive pole of the battery. Because the positive pole tends to accumulate negative charges from the electrons, it also tends to attract positive ions. The ions thus keep charges neutral on both sides and allow the transfer of electrons to go on—at least for a while.
When too many ions have transferred, the battery’s performance begins to decline. Eventually, all the ions that could move have moved, and the battery is discharged. If your battery is rechargeable, you can apply a potential to its electrodes to moves the electrons back to the negative pole the ions will follow suit.
So what is the mysterious quantum physics that governs the flowery arrangements of electron orbitals? That, I am afraid, will have to be the subject of another post.
(*) Note: the attentive reader will have noticed that copper is actually to the left of zinc in the table. However, copper is the more electronegative of the two metals. The rule that electronegativity increases when you move from left to right does have a few exceptions.
While preparing this essay, I benefited from conversations with Piotr Zelenay.