The heaviest element that humans have ever found is called oganesson. Each atom of the stuff packs a whopping 118 protons into its dense center. In contrast, hydrogen—the most abundant element in the universe, something you can find in your body, Earth's oceans and even the atmosphere of Jupiter—has only one. Scientists announced the discovery of oganesson in 2006, when a Russian-American team used a particle accelerator in Dubna, Russia, to fire millions of trillions of calcium ions at a target of heavy atoms. After 1,080 hours of collisions, the investigators had created three atoms of this new superheavy substance.
A few milliseconds later they were gone.
But by carefully accounting for all of the radiation and smaller atoms that the reactions produced, the scientists at the Joint Institute for Nuclear Research in Russia could be fairly sure that they had, for a brief sliver of a moment, created the element. In 2015, after more than a decade of vetting and rechecking, element 118 officially joined the periodic table of the elements, the world's master list of matter. It was named for Yuri Oganessian of the Russian-based institute, who is a pioneer of this research.
But how many more elements are out there? In just the past decade scientists have been pushing the periodic table further and further, adding new atoms that are heavier than ever before. Each one of the fundamental bits of nature on the table is defined by the number of protons packed in its atomic nucleus. At the same time that oganesson became official, researchers also added elements containing 113, 115 and 117 protons per atom to the periodic table. One of us (Düllmann) has been conducting some of the first chemistry experiments on several of the so-called superheavy elements, and Block has been working on the first direct mass measurements and other investigations into some of them. Each new species we find is exciting because it represents an unknown material, a form of matter humans have never encountered before.
We cannot keep them, however. The few atoms we make exist only for brief moments before collapsing or transforming under the strain of too many positively charged protons repelling one another. But scientists suspect that certain yet to be discovered superheavy elements and isotopes (versions of the same element with different numbers of neutrons) might break this teasing pattern of fleeting, tantalizing existence. Some of our envisioned elements may linger for minutes, maybe even years without decaying. If so, they would form a long-sought region on the periodic table called the island of stability. Thanks to special configurations inside their nuclei that grant unusual stability, superheavy elements inhabiting this region may not be just ephemeral creations of the laboratory but could actually stick around. Lately scientists are finding atoms that may represent the shores of this island.
Element 114, for instance, decays somewhat less quickly than some calculations had predicted that an atom jammed with that many protons must. And the half-lives of some of the recently discovered superheavy elements—that is, the time it takes for half of the atoms to decay into another element—get successively longer (although they are still very brief) as their number of neutrons (the chargeless companions of protons in atomic nuclei) increases. These observations fit with predictions—the island is traditionally conjectured to lie in an area of the periodic table where atoms have somewhere around 114 protons and more neutrons than any species created so far. But the discovery of these slightly extended life spans—just a fraction of a fraction of a fraction of a second more—has invigorated a quest that has been a driving force for several generations of heavy element researchers. Now that we have begun to explore the island of stability, we hope to chart its boundaries, determine the location of its heart—where the most stable isotopes lie—and discover how long these atoms can linger.
In recent years scientists have gleaned fascinating insights about these strange denizens of the extreme reaches of the periodic table. Our lab techniques have developed to the point that we can study the chemistry of the superheavy elements we create and determine, for example, whether they would be metals or gases at room temperature. And if we could ever create a kilogram of them, these elements might have completely new—and potentially useful—properties that would distinguish them from all the known materials. Even if the substances we make persist in decaying much too soon for us to ever hold some in our hands, they can help us gain a deeper understanding of chemistry and the fundamental nature of matter.
Exploring the Island
The periodic table is chemistry's ongoing attempt to map that fundamental nature. The chart was developed in the 19th century, most clearly by chemist Dmitri Mendeleev and independently by chemist Julius Lothar Meyer. It lists elements in the order of their atomic number (the number of protons per atom) and lines them up to show similarities in the ways they react with atoms of other elements to form chemical compounds.
Almost as soon as the table was first formulated, chemists wondered how far it extended. The heaviest element found in large quantities in nature is uranium, whose nuclei contain 92 protons. But with each additional proton inside a nucleus, the positive charge grows—and so does something known as the Coulomb force, which repels like charges away from one another. At some point, this push becomes stronger than the attractive “strong interaction” that binds atomic nuclei together, and the nucleus splits apart in a process called fission.
But the stability of any particular element comes down to more than just the number of protons it holds—it depends on the arrangement of protons and neutrons within the atomic nucleus. According to the nuclear shell model developed by Nobel laureates Maria Goeppert Mayer and J. Hans D. Jensen in the late 1940s, both constituent particles fill so-called nuclear shells. Akin to layers within nuclei that can hold specific numbers of protons and neutrons, they are analogous to the electron shells that hold electrons around the nucleus. In both cases, full shells lead to stronger binding, providing extra stability.
Scientists conceived of the shell model when they realized that for specific “magic” numbers of protons and neutrons (2, 8, 20, 28, 50 and 82) nuclei are more stable and harder to rip apart. These magic numbers, it became clear, correspond to full shells. The magic numbers for protons and neutrons in the atoms we know are the same, but there is no guarantee they will continue to match. A nucleus where both proton and neutron shells are full is called doubly magic.
There is much we still do not understand about magic numbers. For instance, what are the magic numbers for nuclei that we have not yet discovered? Some theoretical predictions forecast a doubly magic superheavy nucleus with 114 protons and 184 neutrons. Although we have created element 114 in labs, we have not made a version with 184 neutrons. Yet the prediction of this magic combination, first made in the 1960s, suggested that such an isotope would be so remarkably stable that it would have a half-life approaching the age of Earth. This prediction was the first suggestion of the island of stability—a notion that ignited the field with excitement then and continues to drive us.
We still do not know, however, if 114 and 184 are a true magic combination. Other theoretical frameworks predict, for instance, configurations such as 120 or 126 protons and 172 neutrons. Some of our predictions of future magic numbers owe a debt to Albert Einstein. He explained the surprising observation that the mass of an atom is smaller than the sum of the masses of its constituent protons, neutrons and electrons. His famous formula E = mc2 states that this mass defect reflects the binding energy—the energy that holds the nucleus together. Weighing atoms with different compositions of protons and neutrons thus allows us to identify those configurations that result in a stronger binding—in other words, configurations that represent magic numbers—and to determine how much more stable they are.
Whatever the next magic numbers turn out to be, we think we are starting to encounter the island of stability. Experiments have found that the half-lives of superheavy elements increase with the amount of neutrons, suggesting we are approaching the next magic number of neutrons. This trend is nicely displayed, for example, in the case of element 112 (copernicium, or Cn): compared with Cn 277 (copernicium with 112 protons and 165 neutrons), which lives for only about 0.6 millisecond, Cn 285 (copernicium with eight more neutrons, for a total of 173 neutrons) lives for about 50,000 times longer. The pattern most likely will continue toward the center of the island of stability, although whether any superheavy elements exist that are stable indefinitely is still an open question.
The possibility, though, has triggered a search for these elements in nature. Just because we have not seen them in large quantities, the reasoning goes, does not mean that trace stores of them are not hiding in plain sight. They could have been formed along with other elements heavier than iron in powerful events such as the collision of two dense stars called neutron stars and then seeded throughout the universe. In that case, they could be present in cosmic rays beaming toward us from space or could have survived within rocks on Earth. Scientists have made many searches using different techniques. For instance, element 110 (darmstadtium) is predicted to be quite stable when it has 184 neutrons—a magic number—and might be expected to be chemically similar to platinum, which lies directly above it on the periodic table. Scientists used techniques such as x-ray fluorescence and mass spectrometry to search for the presence of darmstadtium in naturally occurring platinum ores but found no evidence for it in amounts greater than one part in 109.
Researchers have also looked for signs of superheavy elements in cosmic rays—for example, using the Ultra-Heavy Cosmic-Ray Experiment onboard NASA's Long Duration Exposure Facility—but they have found no conclusive evidence. Efforts will certainly continue because such a discovery would be hugely significant. Moreover, new elements could translate to novel materials, each with unique properties that might be useful for technology and other applications.
Forging New Elements
Because we have not yet found any superheavy elements in nature, we must create them ourselves in labs. The task at hand is to enrich the atomic nuclei of “common” elements with even more protons. Up to a point, we can do so by mimicking the process the universe uses to form heavy elements. Nuclei containing too many neutrons are prone to transforming one of their extra neutrons into a proton in the so-called beta decay process, which then produces an element heavier by one atomic number. We can form elements up to fermium (element 100) by bombarding heavy elements with more neutrons. Thus far, though, no fermium nuclei or heavier ones are known to undergo the corresponding beta transformation process, so this pathway ends there.
To go beyond, to create elements such as the elusive oganesson, scientists can bring two nuclei into close enough contact that the strong force kicks in. This force has an extremely short range, meaning that the nuclei must almost touch one another to feel it. To get that close, we have to overcome the repelling force of the positively charged protons, meaning that we must accelerate one of the nuclei to about 10 percent of the speed of light and shoot it at the other. This velocity is just about sufficient to overcome the Coulomb repulsion, whereby the two nuclei make surface contact with each other. But the probability for the two nuclei to touch is extremely small. Furthermore, merging the two initial nuclei into a single combined nucleus gets ever more unlikely the more protons the system contains. Even if such an amalgamated “compound nucleus” does form, it often splits back into lighter fragments practically immediately. Both tendencies—the small probability for a compound nucleus to form in the first place and the large probability for it to split once it is formed—play strongly against the synthesis of ever heavier elements.
Despite the challenges, researchers have achieved exciting successes using this approach. Elements 113, 115, 117 and 118, all made following this pathway, received their official names in 2016. (The International Union of Pure and Applied Chemistry, or IUPAC, is chemistry's version of Guinness World Records and has the power to officially recognize and name new elements.) Element 113 is now called nihonium, honoring Japan, where the experiments to synthesize this element took place. Element 115 is moscovium, honoring the Moscow region, home of the Joint Institute for Nuclear Research, where this element was discovered. And 117 was dubbed tennessine after the state of Tennessee, where Oak Ridge National Laboratory provided the target nuclei of element 97 (berkelium) needed to synthesize the new substance. Oganesson, with its 118 protons, rounds out the new inductees.
Now the heat is on to find element 119, which would add a whole new row to the periodic table. Although several groups, including our own, have gone after these higher numbers, none has succeeded to date, despite investing weeks and months at the world's most powerful accelerators. One obstacle is that the successful route employed up to oganesson—smashing calcium atoms into heavier nuclei—comes to an end at 118 because we do not have sufficient amounts of nuclei with more than 98 protons to use as targets. Scientists are now trying to identify which combinations of known and available elements provide the best chances to generate new species.
Even if these new species decay in a flash, scientists have recently made breakthroughs in their ability to perform experiments during the brief time the elements last to learn about their chemistry and properties—such as whether they behave more like metals or gases at room temperature.
The heaviest element whose chemical properties researchers have studied is flerovium (element 114, or Fl). Flerovium's position on the periodic table, below lead, implies it should be a typical heavy metal. But theories dating back to 1975 suggest it might actually behave more like a noble gas—a very inert gas that rarely interacts with other materials.
The odd expected behavior of flerovium all comes down to how many protons its nuclei contain and, therefore, how highly charged they are. The extreme positive charge of heavy elements' nuclei accelerates their negatively charged electrons to velocities that can reach 80 percent of the speed of light and thereby cause them to revolve around those nuclei in different orbital shapes that have variable spacing than those of the smaller elements. For flerovium, for example, there is a much wider gap in the energy levels of its two outermost electron orbitals than in similar smaller nuclei, such as lead, which sits just above flerovium on the periodic table. In lead's case, forming a chemical bond will more easily provide the energy electrons need to overcome this gap than it will for flerovium. Consequently, flerovium might not undergo chemical reactions as readily as its lighter counterpart. It may therefore resemble other elements that do not like to undergo chemical reactions—the noble gases—more than typical metals such as lead.
It is hard to predict, though, exactly how flerovium will behave. Theories generally agree that it should be more inert than lead, but it is probably more reactive than true noble gases and could form, for example, weak metallic bonds with elements such as gold. Because we have not been able to create it in large enough quantities to observe with human eyes, no one knows what this element would look like in bulk. Some predictions suggest it might appear silvery white or pale gray and take solid form at room temperature.
Flerovium's intriguing properties have inspired scientists to go to extreme lengths to experiment on it, despite the fact that we can produce just single atoms of the stuff per day. Further, even the most long-lived known Fl isotopes have half-lives of only one to two seconds. One of the best facilities we have for producing flerovium is the TransActinide Separator and Chemistry Apparatus (TASCA) at the GSI Helmholtz Center for Heavy Ion Research in Germany. There we shoot a beam of calcium 48 toward a rotating target wheel covered with plutonium 244. When flerovium atoms result, magnets steer them toward a system called the Cryo-Online Multidetector for Physics and Chemistry of Transactinides (COMPACT). This machine consists of two sets of 32-centimeter-long arrays of silicon detectors facing each other, separated by about half a millimeter and forming a narrow, rectangular channel, where a rapidly flowing gas forces flerovium through. The detectors are covered with a very thin gold layer, which allows us to study how the flerovium atoms interact with this metal. The first detector channel is kept at room temperature, but the end of the second channel is cooled with liquid nitrogen to below −160 degrees Celsius because weak bonds—like those exhibited by noble gases—will only be strong enough to bind flerovium atoms at low temperatures but not under warmer circumstances. If flerovium acts more like a metal than a noble gas, it will adsorb on gold on first contact at the warmer beginning of the channel. But noble gases interact too weakly with gold to remain bound at room temperature, so if flerovium behaves more like a noble gas, it will bond in the later part of the channel, if at all.
When our research group used this setup, we observed two atoms that both decayed in the room-temperature detectors, suggesting that flerovium bonded, and then decayed, quickly, more like a metal than a noble gas. Another earlier experiment, though, run by a group at the Paul Scherrer Institute in Switzerland and carried out at the Flerov Laboratory of Nuclear Reactions in Russia, observed three atoms. Although one decayed in the early part of the channel, the two others were found at low temperatures around −90 degrees C. The experimenters interpreted this result as hinting at behavior more like a noble gas. We are currently analyzing data obtained just recently at GSI that we hope will clarify the properties of this exciting element.
Recently one of us (Block) and his team conducted the first laser spectroscopy experiments on a superheavy element—nobelium, or element 102. They were able to produce nobelium atoms at a rate of a few particles per second by bombarding calcium atoms (20 protons) at a target of lead (82 protons). Block and his colleagues then slowed down the resulting nobelium atoms in an argon gas and shot laser pulses at them. If the laser pulses were of the right energy, a nobelium electron would absorb the laser energy and escape from the atom. By varying the frequency of the laser pulses, they were able to precisely measure the energy needed to remove an electron from the nobelium atom. This “ionization energy” is one of the characteristic properties of an element that affects its placement on the periodic table. It determines how likely the element is to react with other elements and form chemical bonds.
We first performed these studies on the isotope No 254 (a version of nobelium with 152 neutrons) and recently extended the experiments to two additional nobelium isotopes, No 252 and No 253, to learn how the different neutron contents shift the energies that the atoms' electrons are able to absorb. The results will tell us about how the size and shape of the nuclei of these isotopes vary—different configurations of the positively charged nuclei will affect the way the negatively charged electrons orbit and behave.
Scientists have also been creating chemical bonds between superheavy elements and more modest atoms to study how the exotic species interact. A recent example is the synthesis of molecules containing seaborgium (element 106). In experiments conducted at the RIKEN Nishina Center for Accelerator Based Science in Japan, Düllmann's group led a team that produced atoms of an isotope of seaborgium with a half-life of about 10 seconds. The scientists then added carbon monoxide to the chamber containing the seaborgium and discovered that the heavy element transformed into a hexacarbonyl compound in which six carbon monoxide molecules bound to the central seaborgium atom.
Düllmann and his colleagues found that in this situation, seaborgium acts much like its lighter and more familiar homologues tungsten and molybdenum, which have the same number of valence electrons. In an approximately two-week-long round-the-clock experiment with seaborgium, the researchers observed seaborgium forming the same type of compound with carbon monoxide molecules as molybdenum and tungsten do, at similar rates. Scientists are now moving to test which of the three elements forms the most stable bonds with carbon monoxide—calculations performed in the late 1990s suggest it will be seaborgium, but more recent and advanced calculations predict the carbon monoxide bond with seaborgium will be weaker than the bond with tungsten.
These are just a few examples of the type of fascinating experiments currently going on with superheavy elements and the many open questions we hope to answer. Although these newest members of the periodic table are admittedly quite exotic, experimental studies provide ever more direct information on how they fit into the same elemental system established by the more common elements that we encounter in our daily life. Whether they be unstable or long-lived—indeed, whether we ever find the heart of the island of stability—superheavy elements have much to teach us about the workings of nature's chemical building blocks.