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(Image: Martin Reznik)

IF IMITATION is the sincerest form of flattery, the periodic table has many true admirers. , and even have been ordered in its image. For chemists, knowing an element’s position in the periodic table, and the company it keeps, is still the most reliable indicator of its properties – and a precious guide in the search for new substances. “It rivals Darwin’s Origin of Species in terms of the impact of bringing order out of chaos,” says of the University of Oxford.

The origins of the periodic table lie in the 19th century, when chemists noticed that patterns began to emerge among the known chemical elements when they were organised by increasing atomic weight. In the 1860s, Dmitri Mendeleev and others began to group the elements in rows and columns to reflect those patterns – and realised gaps in the resulting grids allowed them to predict the existence of elements then unknown.

It was only with the advent of quantum theory in the 20th century that we began to grasp what lies behind these patterns. The periodic table’s rows and blocks roughly correspond to how an atom’s electrons are arranged, in a series of “shells” around the proton-rich nucleus. Electrons fill shells and subdivisions of shells starting with the one closest to the nucleus, which has the lowest energy. The number of electrons in the outermost shell, and its distance from the nucleus and the other shells, are the main factors that determine an element’s chemical behaviour. “Chemical periodicity is a natural property,” says , a philosopher of chemistry at the University of California, Los Angeles.

But that perhaps leads us to some hasty conclusions about the table. “People assume surely it’s been sorted out. It’s not settled – many, many aspects are still up for grabs,” says Scerri. Electron configurations do not always mesh neatly with chemical properties. Properties and patterns we take as given on Earth are very different when we venture into the extreme environment of space. And quite what happens towards the end of the periodic table – and indeed, where this end lies – are questions that remain unanswered. As the following examples show, the periodic table is still very much a work in progress…

Explore our interactive periodic table: “Questions on the table”

Weighty affair

The earliest periodic tables arranged elements by ascending atomic weight – basically, the number of protons and neutrons in an atom’s nucleus. But most atoms come in various isotopes containing different numbers of neutrons. Today’s tables order the elements by atomic number – the unambiguous number of protons.

The atomic weights are still there – but the question is, which is the “correct” one? They used to be displayed as a single number for each element, based on averaging the weights of its natural isotopes according to their relative abundances. But this perpetuates the misconception that this number is some kind of fundamental constant, says of the Reston Stable Isotope Laboratory in Virginia. In reality, the atomic weight of an element such as carbon, say, varies slightly from sample to sample depending on the exact quantities of each isotope.

In 2009 the guardians of the table, the , took action, including hydrogen, lithium, boron, carbon, nitrogen and sulphur, and replacing them with ranges encompassing the isotopic spreads in all known terrestrial samples. Bromine and magnesium followed in May 2013. Nickel, selenium and zinc are probably next up.

Not all elements are so flighty, though. Fluorine, aluminium, sodium, gold and 17 other elements have only one stable isotope, meaning their atomic weight really is a constant of nature. Their weights can stay, then.

Three’s company?

Ordering the periodic table by atomic number makes the position of elements indisputable – except when it doesn’t. Take the case of the two rows of elements floating rather like an afterthought below the main body of the table: the lanthanides and actinides.

Two gaps in the main table, below scandium and yttrium in group 3, mark where these series slot in. The question is, how do they slot in? There are two schools of thought. One goes by electron configurations: scandium and yttrium both have three outer electrons, as do lanthanum and actinium, the elements at the left-hand end of the series, so they are the rightful placeholders. But others point out that chemical properties such as atomic radius and melting point make lutetium and lawrencium at the right end of the rows a better fit. In 2008, in the pages of the Journal of Chemical Education.

“Simmering tensions about the elements’ positions boiled over in 2008”

, says Scerri, and not just for pedagogical clarity. Yttrium can be used to make superconductors, compounds that conduct electricity without resistance, that work at relatively high temperatures. The hunt is on for materials with similar abilities and Scerri thinks lutetium and lawrencium compounds may have been overlooked because they are seen as belonging to a completely unrelated group.

Any resolution will be years away. IUPAC has given Scerri the go-ahead to set up a committee – but only to make the case for why a decision might be needed.

Half empty – or half full?

In the early universe, the two simplest elements, hydrogen and helium, were pretty much all there was. But the advent of more complex stuff has made it difficult to work out where they fit in. “It’s a bit like asking how would you classify dinosaurs along with other animals,” says Scerri.

Hydrogen has one proton surrounded by one electron rattling around in a shell that might hold two. But is that shell half-empty or half-full? Most elements tend to either gain or lose electrons during chemical reactions. Hydrogen swings both ways, sometimes picking up an electron to fill its shell and form compounds like sodium hydride (NaH), and sometimes losing its one electron to form compounds like hydrogen fluoride (HF).

Most periodic tables, , put hydrogen in group 1 with electron-losing metallic elements such as lithium and sodium. But even IUPAC allows that such as fluorine right over in group 17. Most chemists shrug at this ambiguity. “It doesn’t bother me to include hydrogen both in group 1 and in group 17,” says of the University of Helsinki in Finland.

The problem with helium, meanwhile, is that it hardly reacts at all, thanks to its full outer electron shell. In a standard periodic table it sits atop neon with the noble gases that share this characteristic, in group 18. But the fact its outer shell contains only two electrons makes some suspect it would be better off with elements such as beryllium, over in group 2.

That suspicion is increased by calculations indicating that both helium and neon might under certain circumstances react with other elements, but that helium is the more likely to. This goes against the trend of increasing reactivity as you go down that group – a wrinkle that could also be smoothed, , by spiriting helium away to group 2.

When is a metal not a metal?

It was a trend that made the pioneers of the periodic table confident they were on to something: if you sweep diagonally across the periodic table from bottom left to top right, the elements gradually become less metallic. Commonly, the boundary between the two is depicted as a thick line staircasing its way down the table’s right side.

Sadly, it isn’t that simple. “Metal-non-metal status isn’t sacrosanct,” says Edwards. Take hydrogen. We Earthlings encounter it as a distinctly non-metallic transparent gas. But in the cores of hydrogen-rich planets such as Jupiter and Saturn, high pressures and temperatures are thought to make hydrogen a shiny, metallic fluid. Its usual position in the periodic table, above metallic lithium, hints at this. But avowed non-metals such as or are also expected to loosen up under pressure, so their outermost electrons roam free to conduct as they see fit. “The periodic table as you learned it is only the periodic table at ambient conditions,” says of the University of Marburg in Germany.

Hydrogen under pressure might even make a solid metal, a substance with possibly exciting applications as a fuel or room-temperature superconductor – although recent claims to have made it in the lab .

It goes the other way, too. In 2009, a team led by , now at the State University of New York at Stony Brook, used high pressures to turn the shiny group 1 metal sodium into a . In this case, it seems the pressure brings electrons so close that they are forced to occupy spots that minimise subsequent repulsions, instead of roaming free.

Such questionable behaviours highlight that there is little settled in the chemical world, but we shouldn’t panic, says Oganov. “I don’t think we need to revise the periodic table. What we are doing now is making very important comments and corrections to it.”

Einstein’s influence

Einstein’s relativity bends space, time, minds – and the periodic table. By the time you reach gold, with an atomic number 79, the pull of the highly charged nucleus is such that the innermost electrons whizz round at a zippy 80 per cent of the speed of light. This increases their mass, causing them to orbit the nucleus more closely and shield electrons further out from its pull. The outer shells then expand – and the neat connection between how electrons fill up shells and an element’s chemical properties begins to break down.

The knock-on effects on the wavelengths of light that gold absorbs are why it looks so very different from the precious metal directly above it in the periodic table. “You need relativity to actually make gold different from silver,” says Pyykkö. That’s not the only thing. Just last year, at Massey University in Auckland, New Zealand, finally proved something suspected for decades: that mercury’s anomalously low melting point, causing it uniquely among metals to be a liquid at room temperature, .

As ever heavier elements are added to the periodic table, where does this leave things? We’re not quite sure. When the properties of rutherfordium (atomic number 104) and dubnium (105) were found to be out of keeping with hafnium and tantalum immediately above them, questions were asked. But Seaborgium (106) seems stolidly conventional. Element 107 has been dubbed for the way it toes the group line. Experiments with two of the table’s recent additions, copernicium (112) and flerovium (114), have so far painted a mixed picture.

So has Einstein deprived the periodic table of its predictive power? Pyykkö is relaxed about the ructions. “You don’t have a simple mathematical theory underlying the periodic table,” he says. “You have a number of nuts and bolts – one is relativity. When put together, they explain the workings of the periodic table.” , who studies superheavy elements at the GSI Helmholtz Centre for Heavy Ion Research in Darmstadt, Germany, is less sanguine. “The periodic table is still intact – but you can’t predict detailed properties any more,” he says.

Where will it all end?

Early atomic models indicated that, above an atomic number of 103, the repulsive force between positively charged protons would become so great that atoms would just fall apart. Nature generally gives up at 92, with uranium.

But thanks to experiments to fabricate elements beyond uranium in the lab, the heaviest atom now officially recognised is livermorium (116). Elements 117 and 118, already fabricated, are yet to be legitimised. Clearly, where the end of the table lies isn’t that simple. “There certainly is a limit,” says Schädel, “but we don’t know where it is.”

“There certainly is a limit to the periodic table – but we don’t know where it is”

The existence of superheavy elements with atomic numbers above 103 is now explained by theories that say that, just as orbiting electrons are arranged in shells, so too are protons and neutrons within the nucleus. In and around specific “magic” numbers corresponding to full shells, atoms are more stable. But there must still come a point where the fields within highly charged, weighty atoms will become irresistible. One suggestion, from theorist Paul Indelicato of the Pierre and Marie Curie University in Paris, France, and his colleagues in 2011, is that quantum instabilities within these fields .

Stability is only ever relative with the superheavy elements, anyway: those made so far are radioactive and often only survive for fractions of a second before decaying, making them difficult to study. What’s more, most of them can only be produced one atom at a time, meaning chemical properties such as volatility, conductivity, or whether something is a solid, liquid or gas, cease to make much sense. So do they really count as elements at all? “As a chemist, you want to see chemical interactions,” says Schädel.

Several research teams have devised ingenious ways to study the properties of single atoms. One experiment has probed the volatility of copernicium and flerovium by comparing what temperature a single atom sticks to a gold surface with the sticking point of atoms whose volatility is known. Theory also indicates a higher sticking temperature could be evidence for a metallic character. “Chemists and physicists have defined metal for generations but they never had to think about what happens with one atom,” says Schädel. “This is a completely new way of defining a metal.”

Even so, with no realistic prospect of exploiting these superheavy atoms, isn’t this a fool’s errand? Schädel thinks not, and sees such experiments in the 150-year tradition that have made the periodic table the iconic chart of the elements it is. “It’s experiencing terra incognita, going to regions that no one has a glimpsed yet,” he says.

This article appeared in print under the headline “Rogue elements”