The largest and oldest crater on the moon did not form as we initially suspected, a new study suggests. The findings hint that a specific region of the lunar surface could be more scientifically interesting than we thought — with big implications for NASA's upcoming Artemis missions, which are scheduled to land astronauts within this newly realized area of interest as soon as 2027.

The moon formed around 4.46 billion years ago, when an ancient Mars-size protoplanet, known as Theia, collided with Earth, creating a giant cloud of debris that combined into a large spherical satellite orbiting our planet. For the next 200 million years or so, the lunar surface was covered by a hellish magma ocean, as a result of gravitational squeezing from Earth. But as the moon moved farther away from our planet, the molten rock eventually cooled and crystalized, forming an outer rocky crust that has remained largely unchanged ever since — aside from a near-constant bombardment of space rocks.

Shortly after the moon cooled off, around 4.3 billion years ago, an asteroid more than 10 times wider than the one that wiped out the dinosaurs slammed into what is now the satellite's permanent far side. This birthed the South Pole-Aitken (SPA) basin, a giant impact crater around 1,550 miles (2,500 kilometers) across and up to 5 miles (8 km) deep. Researchers believe that this extreme impact may have ejected a mix of radioactive material, dubbed "potassium, rare earth elements and phosphorus" or KREEP, that was left over from the past magma ocean. (The chemical symbol for potassium is K, hence the name KREEP.)

KREEP is believed to have formed during the final cooling phase of the moon, when radioactive elements built up in the space between the moon's newly formed mantle and crust. Studying these elements more closely could help shed light on several mysteries surrounding this period of lunar evolution, including why the crust on the moon's far side is thicker than its visible near side.

Until now, researchers believed that the SPA basin was formed by an asteroid impacting the moon from the south, which would have splashed a layer of KREEP around the crater's northern rim. However, in a new study, published Oct. 8 in the journal Nature, researchers found that the impact actually occurred from the north, meaning that the resulting KREEP is plastered around the southern edge of the crater — right where the crewed Artemis III mission is scheduled to land.

"This means that the Artemis missions will be landing on the down-range rim of the basin — the best place to study the largest and oldest impact basin on the moon," study lead author Jeffrey Andrews-Hanna, a planetary scientist at the University of Arizona, said in a statement. This is "where most of the ejecta, material from deep within the moon's interior, should be piled up," he added.

The researchers became suspicious of the SPA basin's true origins when they compared its shape to other impact craters in the solar system, including Mars' Hellas basin and Pluto's Sputnik basin. All three craters have a similar shape, with one rounded end and another more pointed, making them look somewhat like an avocado or teardrop. The pointed tip of these craters likely represents the direction of impact, enabling researchers to guess which way the SPA basin asteroid hit.

These suspicions were confirmed when the team analyzed data from the Lunar Prospector, a NASA spacecraft that orbited the moon from 1998 to 1999 and measured the radioactivity of surface elements. This revealed a high concentration of thorium — a radioactive element and key component of KREEP — around the crater's southwestern rim.

Collecting KREEP-y samples

The Artemis III mission, which aims to put humans on the moon for the first time since 1972, is currently scheduled for mid-2027, following the completion of the Artemis II mission, which is due to launch at some point before the end of April 2026. The initial mission will send a crewed spacecraft to orbit the moon, similar to the uncrewed Artemis I spacecraft, which successfully circled the moon in late 2022.

The Artemis III astronauts will land near the moon's south pole in one of nine potential sites, most of which lie inside the KREEP splash zone surrounding SPA, according to NASA. If they land in the right place, this could give the team the chance to collect valuable KREEP samples.

However, there is some uncertainty about when Artemis III will actually launch.

Both Artemis II and Artemis III have already been delayed several times. Historically high cuts proposed to NASA's 2026 budget have also led some experts to predict further delays, which may end up handing China the advantage in the race to return humans to the moon.

China has also already acquired the first samples of the moon's far side, which were returned to Earth in June 2024 by the Chang'e 6 spacecraft, from inside the SPA basin. However, despite sharing the valuable lunar rocks with several other countries, NASA has not yet been allowed to anlayze these samples.

Moon quiz: What do you know about our nearest celestial neighbor?

Some atoms are stable, while others seem to fall apart. Lead-208 will probably last forever, while the synthetic isotope technetium-99 exists for just hours. The difference lies in the structure of the atom's nucleus, with certain "magic numbers" of nuclear particles making some isotopes especially resistant to radioactive decay.

So what are these magic numbers, and why are they so special?

The stability of atomic nuclei varies wildly with the number of nuclear particles they contain. Some, like lead-208 and calcium-40, have been around since Earth first formed. Known as primordial isotopes, they will likely survive until the end of time. Others, like oganesson-294 and tennessine-294, are lost to radioactive decay in an instant, with fleeting half-lives of just 0.89 and 0.80 milliseconds, respectively.

This stability seems partly connected to the mass of the atom, with heavier elements proving less stable. But in the 1940s and '50s, scientists observed that many of the lighter elements also had radioactive isotopes; both carbon-14 and potassium-40 undergo radioactive decay slowly and are responsible for much of the planet's background radiation.

Intriguingly, these scientists noticed that very particular numbers of protons and neutrons appeared to result in unusually stable nuclei, and these values became known as magic numbers.

"The magic numbers are 2, 8, 20, 28, 50, 82 and 126," said David Jenkins, a nuclear physicist at the University of York in the U.K. "If you take the lightest one — two protons and two neutrons — that's the nucleus of the helium atom, and we know that's a very stable combination of protons and neutrons."

Related: Why isn't an atom's nucleus round?

Shell game

Helium nuclei, also known as alpha particles, are spontaneously emitted from heavier, unstable atoms as they undergo nuclear decay.

"If you think about it, that's very weird," Jenkins said. "If an atom is going to decay, why doesn't it lose protons or neutrons one at a time? The reason is that the alpha particle is very very stable, and that's related to this idea of magic numbers."

Other magic nuclei include oxygen-16 (eight protons and eight neutrons), calcium-40 (20 protons and 20 neutrons) and lead-208 (82 protons and 126 neutrons), the heaviest stable element known.

To understand these bizarre observations, physicists proposed the "nuclear shell model," which draws parallels with the electronic shells used to explain the chemical behavior of atoms.

"The idea was that protons and neutrons sit in shells, a bit like the electrons in an atom, and nuclear excitations would involve protons and neutrons jumping up and down between those shells," Jenkins explained.

Like their electron analogues, these nuclear shells have fixed energy values known as quantized states, and the system is most stable when these shells are completely filled. The exact reasoning behind this is a complex combination of quantum mechanical factors, but it's thought that the strong force — the fundamental interaction that holds the protons and neutrons together in the nucleus — is higher than expected per particle in completed shells.

Magic numbers are therefore simply the numbers of particles required to fill each of these nuclear shells, with separate levels for protons and neutrons. Individual isotopes can correspondingly be singly magic, with a magic number of either protons or neutrons (for example, the primordial isotope iron-56), or doubly magic, with magic numbers of both protons and neutrons (like oxygen-16 and lead-208).

These doubly magic systems are few and far between, but they possess some intriguing quantum properties, Jenkins said.

"The doubly magic systems have a spherical distribution of matter and charge" — a completely round nucleus, he said. "Most nuclei are deformed and rotate. They have a very different structure."

No one knows how far this model will stretch. Tin-100 — the heaviest doubly magic nucleus, with 50 protons and 50 neutrons — has a half-life of just 1.2 seconds, while unbihexium, the next magic element after lead, has never been synthesized. Therefore, whether this magic stability boost will be enough to allow scientists to add an eighth row to the periodic table remains an open question.

Periodic table of elements quiz: How many elements can you name in 10 minutes?

​​In pursuit of prestige and riches, wealthy people across medieval Europe worked in vain to transmute everyday metals into gold. Today, this process, known as chrysopoeia, is mostly dismissed as an alchemical dream. But is there any science to show that metals can be turned into gold?

In fact, there is — but it would be far from a profitable business, evidence shows.

The idea of transmuting metals to gold goes back to ancient Greece and the philosopher Zosimos of Panopolis. He believed transforming lesser metals into gold was a reflection of the purification and redemption of the soul and the work had a deep spiritual significance. When the concept reemerged in medieval Europe, it was with a purely practical focus — converting a cheap metal into gold was a sure fire route to riches.

"Natural philosophers had this idea of ripening," Umberto Veronesi, an archaeologist and heritage scientist at the NOVA University Lisbon in Portugal, told Live Science. "Base metals were seen as impure stages and would eventually ripen to the purest form of all, which was gold. The only problem was that it would take a very long time for this to happen in the Earth."

Alchemists believed that if they could only create the philosopher's stone — a mythical substance — they would be able to catalyze this ripening process. Metals were thought to contain a mixture of fundamental ingredients: mercury, sulfur and salt. Therefore, by rearranging these components and drawing out any impurities, all metals would ultimately turn to gold, they hypothesized.

"Chrysopoeia was generally consistent with theories of matter and theories of transformation at the time," Veronesi said. "Nobody really doubted that this could be done."

Related: Which is rarer: Gold or diamonds?

The emergence of modern science during the 17th and 18th centuries gradually discredited these ideas, and alchemy was abandoned in favor of the new disciplines of chemistry and physics. However, incredibly, nuclear scientists have held the secrets to this legendary transformation for almost a century.

Today, we know that the identity of an element is determined by the number of protons in its nucleus. Much-coveted gold atoms contain 79 protons, while lead has 82.

"The nucleus is held together by the strong force, and it's very difficult to remove a proton or neutron," said Alexander Kalweit, a physicist working at the Large Hadron Collider at CERN in Switzerland.

However, rearranging these fundamental components of an atom means it's theoretically possible to convert one element into another. "If you have enough energy, you can actually perform such operations," Kalweit said. "When you remove three protons from the lead nucleus, you have created a gold nucleus."

The first successful transmutation of another metal into gold was reported in 1941, when Harvard scientists used a particle accelerator to fire lithium and deuterium nuclei into atoms of mercury, which contains one proton more than gold does. The high-energy particles knocked protons and neutrons from the mercury nuclei, creating three short-lived radioactive isotopes of gold, which quickly decayed because the high-energy nuclei were unstable.

Forty years later, this extraordinary achievement was repeated by Nobel Prize in Chemistry winner Glenn Seaborg at Lawrence Berkeley National Laboratory in California. His team was investigating the fragmentation of bismuth nuclei in relativistic (speed-of-light) collisions and converted several thousand atoms of the element into gold by bombarding the sample with carbon and neon nuclei in a particle accelerator.

Today, research teams at particle accelerators around the world continue to report the production of gold as a by-product from their experiments. At the Large Hadron Collider, Kalweit's team is investigating the collisions of lead ions at close to the speed of light.

"In the head-on collisions, we essentially liberate the quarks that are inside the protons and neutrons, and they, for a short time, form a state of matter that existed a few microseconds after the Big Bang in the early universe," he explained. "It's the so-called quark gluon plasma."

These head-on collisions are so intense that the protons and neutrons are completely destroyed. But lower-energy near-miss interactions — where the particles are extremely close but not touching — generate a powerful electromagnetic field that knocks protons out of the lead nuclei. The result: The team detected around 29-trillionths of a gram of gold during a three-year experimental run.

However, despite having achieved the alchemist's dream, it's unlikely that nuclear physicists will ever turn a profit by synthesizing gold in a particle accelerator. The expense of building and running a facility like the Large Hadron Collider is astronomical compared with the value of the volume of gold produced; it's estimated that Seaborg's experiments in the 1980s cost around a trillion times the price of the gold they produced. Plus, the rarity of interesting interactions means researchers must wade through billions of data points to even identify the transformed atoms.

"Since the 1940s, there are many experiments which have produced gold," Kalweit said. "But what is common to all of them is that none of them is even remotely close to being profitable."

Periodic table of elements quiz: How many elements can you name in 10 minutes?

Scientists may have finally worked out where gold came from in the early universe.

In a new study, researchers revealed that powerful flares originating from neutron stars with a strong magnetic field, called "magnetars," may have started forging gold not long after the Big Bang — significantly earlier than previously thought possible. The researchers described their findings in a study published Tuesday (April 29) in The Astrophysical Journal Letters.

Scientists have long been puzzled by the origins of the universe's vast amounts of gold. Researchers already knew that mergers of collapsed stars and black holes discharge heavy metals, but in 2017, for the first time ever, astronomers observed the merger of two superdense stellar corpses known as neutron stars. The cataclysmic collision, which happened 130 million light-years away, emitted a flash of light that contained signatures of heavy metals, including platinum and staggering amounts of gold.

But while the 2017 event accounted for some of the universe's gold abundance, it couldn't explain how gold and heavy metals formed in the universe's early days, because not enough time had passed for neutron star mergers to occur.

Now, scientists think they can finally explain how gold and other heavy elements were first created and distributed in space.

"It's answering one of the questions of the century," study co-author Eric Burns, an assistant professor of physics and astronomy at Louisiana State University, said in a NASA statement.

Related: Hubble watches neutron stars collide and explode to create black hole and 'birth atoms'

Forged by 'extreme explosions'

Magnetars have existed since the early days of the universe, and the study's authors estimate that these structures may have contributed up to 10% of all elements heavier than iron in the Milky Way, according to the statement.

The researchers used 20-year-old data from NASA and European Space Agency (ESA) telescopes to find the universe's hidden source of gold and heavy metals. They narrowed their search to magnetars based on the results of a 2024 study, which found that magnetar giant flares — bursts of radiation released during "starquakes" — can eject material, including heavy metals, from the crust of neutron stars and into space.

The last magnetar giant flare observed from Earth was in 2004. Scientists at the time noticed a small gamma ray signal from the flare, "but nobody had any conception of what it could be," Burns said.

It turns out, this small signal mirrors signals that scientists would expect to see if a magnetar created and threw out heavy metals in a giant flare.

Magnetar giant flares produce an enormous amount of high-energy radiation, which could be the key to forging gold and other elements heavier than iron, according to the researchers. Specifically, the authors of the new study think that the extremely high density of neutrons in a giant flare could transform light atomic nuclei into much heavier ones, triggering multiple nuclear decay reactions in a single atom at once.

Atoms carry protons and neutrons, which determine an element's identity and mass, respectively. Hydrogen is the simplest atom on the periodic table, because it has only one proton. Helium, the second-simplest element, has two protons; lithium has three, and so on.

Under certain conditions, atoms can absorb an extra neutron, which increases the mass of the atom, destabilizing it and sparking a nuclear decay reaction that converts this neutron into a proton. When that happens, the atom that absorbed the neutron has an extra proton, which changes its identity and moves it up the periodic table. Hydrogen becomes helium, helium becomes lithium, and so forth.

Magnetic giant flares host a turbocharged version of this process, because the huge density of neutrons can cause atoms to absorb several of them at once, according to the researchers. Thus, a relatively light atom may suddenly transform into a much heavier one, leading to the rapid formation of heavy metals, including gold.

"It [is] very cool to think about how some of the stuff in my phone or my laptop was forged in this extreme explosion [over] the course of our galaxy's history," study lead author Anirudh Patel, a doctoral student in astrophysics at Columbia University in New York, said in the statement.

The next step for the researchers is to look for further clues in older magnetar giant flare data. NASA's Compton Spectrometer and Imager (COSI) mission will also follow up on the results when it is launched, which is expected in 2027. COSI is a wide-field gamma ray telescope that will study energetic phenomena in the cosmos, including magnetar giant flares.

For the first time, scientists have revealed crucial properties of the mysterious, radioactive substance promethium — nearly eight decades after the elusive rare earth element was discovered.

Promethium is one of the 15 lanthanide elements at the bottom of the periodic table. Also known as the rare earths, these metals exhibit a number of useful properties, including strong magnetism and unusual optical characteristics, making them particularly important in modern electronic devices. 

"They are used in lasers; they are part of the screens of your smartphone. They are also used in very strong magnets in wind turbines and electric vehicles," Ilja Popovs, a research and development staff member at Oak Ridge National Laboratory (ORNL) and co-author of a new study published in the journal Nature, told Live Science. 

'Scarce and difficult to study'

Promethium itself, which was discovered by ORNL scientists in 1945, has a few minor applications in atomic batteries and cancer diagnostics. But scientists have a very limited understanding of the element's chemistry, precluding more widespread uses. 

Studying the radioactive element has posed a decades-long challenge, partly due to the difficulty of securing a suitable sample, team member Alexander Ivanov, also a research and development scientist at ORNL, told Live Science.

"Promethium doesn't have a stable isotope — they're all radioactive, meaning that they are decaying [into other elements] with time," Ivanov said. "You get this element through a fission process, so it's scarce and difficult to study."

ORNL is the U.S.' only producer of promethium-147, an isotope of the element with a radioactive half-life of 2.6 years. Using a method developed last year, the researchers separated this isotope from nuclear reactor waste streams, creating the purest possible sample for study.

Related: Atoms squished closer together than ever before, revealing seemingly impossible quantum effects

Then, the team combined this sample with a ligand — a molecule specially designed to trap metal atoms — to form a stable complex in water. The coordinating molecule, known as PyDGA, formed nine promethium-oxygen bonds, giving researchers the first-ever opportunity to analyze the bonding properties of a promethium complex.

However, the analysis itself was no trivial matter. 

"Because promethium is radioactive, once it's decaying, it's getting transmuted into the adjacent element, which is samarium," Ivanov said. "So you will have a tiny amount of contamination in the form of samarium." 

'The last puzle piece'

The team therefore used an extremely specialized, element-specific technique called synchrotron-based X-ray absorption spectroscopy. High-energy photons generated by a particle accelerator bombarded the promethium complex to build a picture of the positions of atoms and the lengths of bonds. Subtle differences in the metal-oxygen bond lengths then allowed the team to focus on the key promethium-oxygen bond, discounting any contaminating samarium.

Crucially, this information enabled a comparison of promethium's properties with other rare earth complexes for the first time. 

"Promethium was the last puzzle piece among those elements," Popovs said. The ligand provided a way to have a stable complex for all of the lanthanides — the same element ratios and the same kind of geometry. That allowed the team to "study the fundamental physical chemical properties of these complexes across the whole series," Popovs explained.

Lanthanides are naturally found as mixtures of elements, so understanding periodic trends such as bond lengths and complex-forming behaviors helps scientists develop new and more efficient methods to separate these valuable metals. 

Now, the ORNL team is studying promethium in water to build a clearer picture of the coordination environment and chemical behavior of this unusual element. 

"Hopefully, the fundamental insights that we're providing will inform other scientists how to design better separation technologies and can perhaps spur more interest in studying it for other applications," Popovs said.

​​Bismuth is an unusual element that we don't encounter much in everyday life. But this pretty, iridescent metal, found near the bottom of the periodic table, exhibits some extraordinary properties. Magnetic levitation — bismuth's ability to seemingly float between two magnets — is perhaps one of the most interesting. The repulsion between bismuth and the magnets is so strong, it causes the metal to levitate.

But why is bismuth so strongly repelled from magnets?

According to Eric Riesel, a magnetic materials chemist at MIT, the answer comes down to the type of magnetism exhibited by bismuth. Every material has magnetic properties, determined by a quantum property of the element's electrons known as spin. But, this spin can only point in two directions — up or down — and the combination of all the spins in a material define exactly what type of magnetism the element will exhibit.

"Most people are familiar with ferromagnets (permanent magnets) like iron, where the spins are all aligned with each other, but there are also anti-ferromagnets where the spins are pointed in opposite directions to each other," Riesel told Live Science.

However, there's also another pair of magnetic categories: paramagnetism and diamagnetism. "In paramagnets, when you apply a magnetic field, spins in that material will align with the field in proportion to its strength," he said. "Diamagnets apply a force in the opposite direction to the field, repelling it."

Related: Is it possible to reach absolute zero?

Bismuth is an example of a diamagnetic material, yet this is not the behavior we would expect from the element's electron configuration. The type of magnetism exhibited by a material depends on the arrangement of electrons and their corresponding spins. Electrons circle the nucleus in defined layers called shells, which are further subdivided into levels called the s, d, p and f orbitals.

Typically, diamagnetic materials have a closed shell structure. This means a particular group of orbitals are completely full and the electrons have been forced to pair, with one pointing up and the other down — essentially canceling out the spins. Conversely, paramagnetic materials usually have partially filled orbitals, meaning the electrons are unpaired and can align their spins in the same direction.

Bismuth is in Group 15 of the periodic table. The s, d and f orbitals are all full, but the p orbitals contain three out of a possible six electrons. So bismuth has partially filled orbitals and should behave as a paramagnet. However, its position in row six of the periodic table means bismuth also possesses some unusual heavy-atom properties.

"Chemical elements found after the f-block in the periodic table have their outermost electrons orbiting the nucleus at speeds that are significant fractions of the speed of light," said Ira Martyniak, also a magnetic materials chemist at MIT. "The direct relativistic effect makes the 6s and 6p orbitals contract and reside closer to the nucleus, which gives rise to anomalous physical and chemical characteristics."

These relativistic effects are responsible for many of bismuth's surprising properties, such as its unconventional superconductivity, its very low melting point (520.7 degrees Fahrenheit, or 271.5 degrees Celsius) and the unusual shape of its crystals. The unexpected diamagnetism is no exception.

"Even though bismuth has the unpaired electrons in its 6p orbital, because of relativistic contraction of the 6s and 6p levels, the paramagnetism stemming from the 6p electrons is suppressed and the behavior of bismuth is largely dominated by the closed shells and large size of the atom, leading to strong diamagnetism," Martyniak told Live Science.

Diamagnetic materials have lots of valuable applications, including electromagnetic induction in copper coils (used to generate electricity) and the aluminum tracks of high-speed maglev trains. Bismuth itself is too heavy to be a practical material for general use, but its potent diamagnetism means it is now a common component in superconductors and quantum computing.

Scientists have discovered evidence of a massive star from the early universe that does not fit with our current understanding of the cosmos. The ancient stellar oddball, which researchers have dubbed the "Barbenheimer Star," likely had a mix of elements in its core that has never been seen before — then, it died a seemingly impossible death while birthing an equally puzzling star in its place, a new study shows. (The name Barbenheimer is a reference to the contrasting films "Barbie" and "Oppenheimer" releasing on the same day last year.)

Researchers uncovered traces of the Barbenheimer Star after taking a closer look at J0931+0038, a distant red giant star. J0931 was first discovered in 1999 by the Sloan Digital Sky Survey (SDSS) — one of the largest and most detailed astronomical databases of the night sky — but had not been properly analyzed until now. 

In a new study uploaded to the preprint server arXiv on Jan. 4, researchers turned the SDSS telescopes in New Mexico back toward J0931 and captured a detailed spectrum of the star's light, which was later verified by follow-up observations from the Giant Magellan Telescope in Chile. These spectra revealed that J0931 seemingly had an extremely odd metallicity, or chemical composition, with an unusually high concentration of heavy elements. (These results have not yet been peer-reviewed.)

Using the newly acquired data, the research team pieced together how J0931 formed via a process known as stellar archaeology. This revealed that the star was birthed from the supernova remnant of an even larger star — between 50 and 80 times more massive than the sun — that dates back as far as 13 billion years ago, only around 700 million years after the Big Bang.

The metallicity of the parent star (Barbenheimer) was likely equally as weird as that of J0931 before it blew up, which would have been completely different from other known stars in the early universe.

"We've never seen anything like this," study lead author Alex Ji, an astrophysicist at the University of Chicago, said in a statement. "Whatever happened back then, it must have been amazing."

Related: 1st evidence of nuclear fission in stars hints at elements 'never produced on Earth'

J0931's metallicity was strange for three reasons. First, the star had unusually low levels of lighter elements such as magnesium, sodium and aluminium, which are normally more abundant in stars. Second, it had an unusually high amount of midweight elements such as iron, nickel and zinc. And finally, it had an "overabundance" of heavier elements like strontium and palladium, according to the researchers.

"We sometimes see one of these features at a time, but we've never before seen all of them in the same star," study co-author Jennifer Johnson, an astronomer at The Ohio State University, said in the statement.

Most stars have the reverse metallicity of J0931: They have higher levels of lighter elements and lower levels of midweight and heavier elements. This is because stars are made predominantly of hydrogen and helium, which fuse together in the stars' cores to create heavier elements. These new elements, which are much less abundant, eventually fuse into heavier and heavier elements. 

It is therefore hard to explain why J0931 has such and abundance of heavy elements because it doesn't seem to have a high enough concentration of lighter elements to have created them. 

"Amazingly, no existing model of element formation can explain what we see," said study co-author Sanjana Curtis, an astronomer at the University of California, Berkeley. It "almost seems self-contradictory," she said.

J0931's unusual metallicity would have partially been inherited from the ingredients that the Barbenheimer Star spit out when it exploded. This means that the parent star would likely have had a similarly inverted metallicity. This is even stranger, because in the early universe, stars shouldn't have existed long enough to have created such high concentrations of heavy elements, the team said.  

But what's even stranger is that the Barbenheimer Star should have never gone supernova, the researchers wrote. In theory, a star with Barbenheimer's predicted mass should have collapsed into a black hole rather than exploding outward. At the moment, the study team cannot explain why this collapse didn't happen. 

The only way for scientists to learn more about the Barbenhaimer Star and its bizarre composition is to search for other similar stellar oddballs from the early universe to uncover more pieces of this cosmic puzzle.

"The universe directed this movie, we are just the camera crew," study co-author Keith Hawkins, an astronomer at the University of Texas at Austin, said in the statement. "We don't yet know how the story will end."

The universe is flooded with billions of chemicals, each a tiny pinprick of potential. And we've only identified 1% of them. Scientists believe undiscovered chemical compounds could help remove greenhouse gases, or trigger a medical breakthrough much like penicillin did.

But let’s just get this out there first: it's not that chemists aren't curious. Since Russian chemist Dmitri Mendeleev invented the periodic table of elements in 1869, which is basically a chemist's box of Lego, scientists have been discovering the chemicals that helped define the modern world. We needed nuclear fusion (firing atoms at each other at the speed of light) to make the last handful of elements. Element 117, tennessine, was synthesised in 2010 in this way.

But to understand the full scale of the chemical universe, you need to understand chemical compounds too. Some occur naturally — water, of course, is made of hydrogen and oxygen. Others, such as nylon, were discovered in lab experiments and are manufactured in factories.

Elements are made of one type of atom, and atoms are made of even tinier particles including electrons and protons. All chemical compounds are made of two or more atoms. Although it's possible there are undiscovered elements left to find, it's unlikely. So, how many chemical compounds can we make with the 118 different sorts of element Lego blocks we currently know?

Big numbers

We can start by making all the two-atom compounds. There are lots of these: N2 (nitrogen) and O2 (oxygen) together make up 99% of our air. It would probably take a chemist about a year to make one compound and there are 6,903 two-atom compounds in theory. So that's a village of chemists working a year just to make every possible two-atom compound.

There about 1.6 million three-atom compounds like H₂0 (water) and C0₂ (carbon dioxide), which is the population of Birmingham and Edinburgh combined. Once we reach four- and five-atom compounds, we would need everyone on Earth to make three compounds each. And to make all these chemical compounds, we'd also need to recycle all the materials in the universe several times over.

But this is a simplification, of course. Things such as the structure of a compound and its stability can make it more complex and difficult to make.

The biggest chemical compound that has been made so far was made in 2009 and has nearly 3 million atoms. We're not sure what it does yet, but similar compounds are used to protect cancer drugs in the body until they get to the right place.

But wait, chemistry has rules!

Surely not all those compounds are possible?

It's true there are rules — but they are kind of bendy, which creates more possibilities for chemical compounds.

Even the solitary "noble gases" (including neon, argon and xenon and helium), which tend to not bind with anything, sometimes form compounds. Argon hydride, ArH+ does not exist naturally on Earth but has been found in space. Scientists have been able to make synthetic versions in laboratories that replicate deep space conditions. So, if you include extreme environments in your calculations, the number of possible compounds increases.

Carbon normally likes being attached to between one and four other atoms, but very occasionally, for short periods of time, five is possible. Imagine a bus with a maximum capacity of four. The bus is at the stop, and people are getting on and off; while people are moving, briefly, you can have more than four people actually on the bus.

Some chemists spend their entire careers trying to make compounds that, according to the chemistry rulebook, shouldn't exist. Sometimes they are successful.

Another question scientists have to grapple with is whether the compound they want can only exist in space or extreme environments — think of the immense heat and pressure found at hydrothermal vents, which are like geysers but on the ocean floor.

How scientists search for new compounds

Often the answer is to search for compounds that are related to ones that are already known. There are two main ways to do this. One is taking a known compound and changing it a bit — by adding, deleting or swapping some atoms. Another is taking a known chemical reaction and using new starting materials. This is when the method of creation is the same but the products may be quite different. Both of these methods are ways of searching for known unknowns.

Coming back to Lego, it's like making a house, then a slightly different house, or buying new bricks and adding a second storey. A lot of chemists spend their careers exploring one of these chemical houses.

But how would we search for truly new chemistry — that is, unknown unknowns?

One way chemists learn about new compounds is to look at the natural world. Penicillin was found this way in 1928, when Alexander Fleming observed that mould in his petri dishes prevented the growth of bacteria.

Over a decade later, in 1939, Howard Florey worked out how to grow penicillin in useful amounts, still using mould. But it took even longer, until 1945, for Dorothy Crowfoot Hodgkin to identify penicillin's chemical structure.

That's important because part of penicillin's structure contains atoms arranged in a square, which is an unusual chemical arrangement that few chemists would guess, and is difficult to make. Understanding penicillin's structure meant we knew what it looked like and could search for its chemical cousins. If you're allergic to penicillin and have needed an alternative antibiotic, you have Crowfoot Hodgkin to thank.

Nowadays, it's a lot easier to determine the structure of new compounds. The X-ray technique that Crowfoot Hodgkin invented on her way to identifying penicillin's structure is still used worldwide to study compounds. And the same MRI technique that hospitals use to diagnose disease can also be used on chemical compounds to work out their structure.

But even if a chemist guessed a completely new structure unrelated to any compound known on Earth, they'd still have to make it, which is the hard part. Figuring out that a chemical compound could exist does not tell you how it's structured or what conditions you need to make it.

For many useful compounds, like penicillin, it's easier and cheaper to "grow" and extract them from moulds, plants or insects. Thus the scientists searching for new chemistry still often look for inspiration in the tiniest corners of the world around us.

This edited article is republished from The Conversation under a Creative Commons license. Read the original article.

The excerpt below is taken from "Elemental: How Five Elements Changed Earth’s Past and Will Shape Our Future" (Princeton University Press, 2023) by Stephen Porder. It looks at how one of the biggest events in Earth's history came to be: plants colonizing the continents.

Plants are so ubiquitous on land that it's hard to imagine their absence, but for almost 90% of Earth's history, there was no life on land, or at the very least no plants. Land plants emerged a little more than 400 million years ago, which compared with the 4-billion-year history of life on Earth, is comparatively recent. 

This evolutionary leap allowed them to become (arguably) only the second group of organisms to radically change the world, a full 2 billion years after the first, cyanobacteria, oxygenated the planet. Their world-changing predecessors precipitated a Great Oxidation Event that was likely the biggest environmental disaster in history — but also set the stage for all multicellular life on Earth. Land plants did not have quite this big an impact, but theirs was greater than any other group of organisms in the intervening two billion years. 

For reasons I'll explore below, plants' evolutionary innovations are, in some ways, best understood through their connection to their cyanobacterial predecessors, and to the next group of world-changing organisms to evolve — humans. That connection lies in Life's Formula, the five elements that make up all living things: hydrogen, oxygen, carbon, nitrogen and phosphorus.

Let's set the stage for the story of plants by considering the world into which they emerged. The oceans of 400 million years ago were not radically different from those that cover 70% of Earth's surface today. Earth was much warmer, however, with a tropical climate from pole to pole. All the major types of life, including vertebrate and invertebrate animals of all shapes and sizes, roamed the seas. The continents peeking out from the watery surface were about their modern sizes, though not in their modern locations. Crucially, the chemistry of the ocean was similar to today, and understanding that chemistry helps explain how truly extraordinary, and world changing, the move to land was. 

How did organisms live in that ancient ocean? As today, ocean food chains were built on the consumption of oxygen-producing photosynthetic organisms like cyanobacteria and plankton. The cellular machinery of photosynthesis in these single-celled organisms was sustained by nitrogen inputs from cyanobacteria and other microorganisms that could "fix" nitrogen from the limitless supply in the air. 

Water (hydrogen plus oxygen) and nitrogen are three elements in Life's Formula, the elements that all living things share in very similar ratios. Sunlight, abundant water and "fixed" nitrogen fueled the acquisition of a forth — carbon — through photosynthesis. Despite the fact that cyanobacteria can tap into a virtually limitless source of nitrogen from the air, we think nitrogen placed a key limitation on how much life existed in the ancient ocean (that constraint remains today). It is a bit of a puzzle as to why this would be the case. Nitrogen fixation may give cyanobacteria a leg up on photosynthetic organisms that cannot pull off this remarkable bit of biological alchemy. But once a cyanobacteria cell dies and decomposes, the nitrogen it captured should become available to other organisms. Recycling is the norm in nature — once a scarce nutrient enters a system, it tends to stay there, fiercely sought after by all concerned. So why, when cyanobacteria could tap into a virtually limitless bank account of nitrogen in the air, did nitrogen remain relatively scarce in the ocean? Why didn't the cyanobacteria cause it to accumulate until it was no longer in short supply?

Related: Tropical rainforests could get too hot for photosynthesis and die if climate crisis continues, scientists warn

This puzzle has preoccupied scientists in my field for decades, and like many good puzzles there is no single, clear answer. Nitrogen losses are clearly important, but I want to focus on another among the many reasons people have come up with: that the proliferation of cyanobacteria specifically, and photosynthetic organisms in the ocean in general, was limited by another element in Life's Formula. 

The most abundant two elements in Life's Formula are hydrogen and oxygen. Living in the ocean, cyanobacteria had plenty of access to these. Photosynthesis uses sunlight and water efficiently to capture carbon, of which there is no shortage in the ocean. Research from as early as the 1950s showed convincingly that enough CO2 gas dissolves into the ocean that it rarely, if ever, is a constraint to growth. Photosynthetic machinery requires lots of nitrogen, but cyanobacteria can fix nitrogen, which dissolves in ocean water since it is so abundant in the air. And then … there's phosphorous. 

It turns out that organisms that can fix nitrogen tend to have high demands for other atoms — particularly phosphorus, but also iron and molybdenum. The latter two are important components of the biological machine (the nitrogenase enzyme) that carries out nitrogen fixation. Phosphorus, iron, and molybdenum, unlike nitrogen, are virtually absent from the air. They are made available to organisms by the chemical breakdown of rock, and thus, with an admitted lack of linguistic imagination, scientists call them "rock-derived." We now think that these rock-derived elements limit the growth of cyanobacteria and other nitrogen-fixing organisms in the oceans. Thus, while life might have been proximately limited by the amount of nitrogen, the amount of nitrogen those organisms could capture was ultimately limited by the supply of elements derived from the weathering of rocks. 

Imagine yourself as a single-celled, photosynthetic organism floating in the middle of the ocean 400 million years ago, more than 1,000 miles from land. If you're at the surface, there is plenty of sunlight available to drive photosynthesis. There are plenty of water molecules to split using the energy from the sun. If you're a nitrogen fixer, like cyanobacteria, you can build the machinery to capture nitrogen gas that is dissolved in the water. But where do you get the elements — the rock-derived phosphorus, iron, and others — needed to build that machinery? Not from the weathering of rocks at the ocean bottom — those are miles down — and even if you managed to get down there, there wouldn't be any light to fuel photosynthesis. As a single-celled organism in the upper ocean, you would just have to wait and hope that those elements come to you.

But if you're an unlucky single-celled organism, you live in a vast ocean desert. These places have very little life, despite being replete with sunlight and CO2, because they lack the other elements of Life's Formula. The only source of rock-derived phosphorus, for example, is the transport of material from the continents — a slow trickle of dirt from rivers and dust falling on the ocean surface. Floating in the middle of the Paleo-Pacific Ocean, you are at the mercy of the currents. There are no rocks for miles: up, down, or sideways. There is nothing you can do to increase your access to rock-derived elements. No way to access the fifth-most abundant element in your cells — phosphorus — and the other atoms derived from the breakdown of rocks. No way, that is, except to evolve and move to the source: land.

As with the cyanobacterial revolution that oxygenated the planet, the evolutionary innovations that allowed plants to complete the slow march landward revolved around access to the elements in Life's Formula. A first, and critically important, step was to bring the photosynthetic machinery from the ocean with them. The chloroplasts in plant leaves — the place where photosynthesis occurs — have their own DNA. It's the DNA of photosynthetic ocean bacteria that, long ago, merged into plant cells. Chloroplasts are thus an example of endosymbiosis — an organism within an organism. As a result of this endosymbiosis, the chemical reaction of plant photosynthesis is the same as cyanobacteria photosynthesis. It uses the same machinery. That is why land plants pump out oxygen during photosynthesis in the same way cyanobacteria do.

Living in the ocean meant using water for photosynthesis wasn't a problem. But on land, the need for water means a constant struggle to stay hydrated. The struggle is encapsulated by Life's Formula, which starts with hydrogen and oxygen. Because land plants inherited their photosynthetic machinery from their ocean-dwelling, single-celled ancestors, they use the same hyper-efficient, water-reliant photosynthesis. They split water using the energy from sunlight, capture CO2, and produce sugars to build their cells (and oxygen, by evolutionary accident). But every moment they open their leaves tiny pores to let CO2 diffuse in from the air they lose scarce water out through the same conduit. This is a scarcity ocean-dwellers don't have to deal with.

The evolutionary solution to this scarcity was the development of water saving mechanisms: leaf waxes, extensive root networks, and symbiosis with fungi that explored every nook and cranny of the soils. These innovations gave access to water, and as roots and fungi attacked the rocks below, they liberated phosphorous as well. These rocks were far out of reach for plants' ocean-dwelling predecessors, but right beneath their "feet" on land. By chemically and physically attacking the rocks upon which they grew, plants and their fungal partners became the world's first, and most efficient, miners, and gained greater access to the key elements in Life's Formula.

By colonizing the continents and moving to the source of the elements whose availability constrained their ocean-dwelling ancestors, land plants set themselves up to become the second great world-changers. To understand how, we have to move from understanding the paleo-ocean to understanding the paleo atmosphere. As today, nitrogen (as N2 gas, two nitrogen atoms bound so tightly together they are virtually inert) and oxygen (as O2 gas, two oxygen atoms bound together loosely enough to be very reactive) made up the vast majority of the air. But the best available evidence suggests CO2 levels may have been ten times higher than today, and the heat trapped by all that CO2 meant the world was very hot, probably about 10 degrees Fahrenheit (5.5 degrees Celsius) hotter than today. This may not sound like a lot, but such a world was hot enough to have no ice at either pole, the northern one covered by bathtub-temperature ocean and the southern by the supercontinent Gondwana.

Land plants made three key innovations. First, they found a new way to capture sunlight and thus carbon. In this case the innovation wasn't a new biochemical reaction but the movement of this reaction to a new place. Second, they evolved a way to withstand water scarcity on land by building root networks and partnering with fungi (among other things). Finally, they became miners, digging for critical rock derived nutrients that were, and remain, scarce in the ocean. Their innovations in getting water and nutrients allowed their wild proliferation. Proto-forests spread across much of the supercontinent that spanned from equator to pole. But, as with cyanobacteria, the story of plants also shows how unprecedented access to life's essential elements can have consequences. Once again, innovation and proliferation ended with catastrophe.

The catastrophe came about because the elements in Life's Formula are also contained in the greenhouse gases that regulate Earth's climate. As today, 400 million years ago the main gas keeping the planet warm was CO2. When plants evolved, they pulled CO2 from the air to build their tissues, and when those tissues died, some of that carbon got stuck in soils. Withdrawal #1 from the bank of CO2 in the air. Plants also accelerated the dissolution of minerals on land, which had the net effect of removing CO2 from the air and storing it on the ocean floor as limestone. Withdrawal #2. Finally, geologic conditions allowed the growth and repeated flooding of the vast lowland swamp forests that emerged during what is, not coincidentally, known as the Carboniferous Period. When plants growing in those swamps died, their remains were protected from decomposition. Their burial, over millions of years, represented yet another net transfer of CO2 out of the air. Withdrawal #3. All else being equal, you can't increase the rate at which you withdraw from a bank account without having that account go down. With the triple whammy of withdrawals that land plants imposed, the amount of CO2 in the air began to fall.

Eventually, plants' innovations pulled enough CO2 out of the air that the greenhouse effect began to weaken. The pan-tropical Earth, which had supported great forests across most of its land, began to cool. It is unclear how long the process took before Earth chilled enough to have ice ages. But by 300 million years ago, roughly 100 million years after plants got going in earnest on land, Earth had cooled enough that the vast tropical forests were gone from most of the planet. They were frozen by their own success. An environmental disaster spurred by new access to the elements in Life's Formula, subsequent proliferation, and collateral consequences.

The process driven by plants was slow: a drip, drip, drip out of the bank account of CO2 in the air and a transfer of that carbon below ground. Some of that carbon was gradually compressed, concentrated, and turned into coal. Then, 300 million years after those tropical trees succumbed to environmental changes of their own making, the next world-changing organism, humans, discovered that carbon-rich bank account. 

We began burning this stored carbon at a rate never before seen in the history of our planet. We used the energy that burning produced to build dams and capture water, allowing us and our crops to stay hydrated on land. We used that energy to industrially fix nitrogen and mine phosphorus to fertilize our now-irrigated farms. And we too, are changing the world, even faster than our predecessors. But like them, our success, and environmental peril, is tied inextricably linked to the elements in Life's Formula.

Text from ELEMENTAL by Stephen Porder. Copyright © 2023 by Princeton University Press. Reprinted by permission of Princeton University Press.

The James Webb Space Telescope (JWST) would be able to spot the signs of our civilization on Earth if it was spying on us from another star system in the Milky Way, a new study shows. The finding raises hopes that the state-of-the-art spacecraft could detect alien civilizations as it stares out toward distant worlds in our galaxy.

Since launching in late 2021, JWST has been predominantly peering out into the deepest reaches of the cosmos in search of clues about how the early universe formed. But one of the telescope's secondary objectives is to analyze the atmospheres of nearby exoplanets, or planets beyond the solar system, to look for gases produced by biological life, known as biosignatures, and chemicals produced by advanced alien civilizations, known as technosignatures.

Related: Are aliens real?

But despite being the most advanced telescope currently in operation, it is still unclear how well JWST will be able to spot the tell-tale signs of intelligent life. To answer this question, researchers decided to test whether the space telescope could successfully detect intelligent life from the only planet in the universe that is known to be both habitable and currently inhabited — Earth.

In the new study, uploaded to the pre-print server arXiv on Aug. 28, researchers took a spectrum of Earth's atmosphere and deliberately decreased the quality of the data to mimic how it would look to an observer dozens of light-years away. The team then used a computer model, which replicated JWST's sensor capabilities, to see if the spacecraft could detect the key biosignatures and technosignatures from the dataset, such as methane and oxygen, produced by biological life, and nitrogen dioxide and chlorofluorocarbons (CFCs), which are produced by humans.

The results, which have not yet been peer-reviewed, show that JWST could likely detect all the key markers of non-intelligent and intelligent life in our planet's atmosphere.

Related: There may be hundreds of millions of habitable planets in the Milky Way, new study suggests

The researchers noted that the quality of the altered dataset is roughly equivalent to JWST observations of planets from TRAPPIST-1 — a star system containing seven exoplanets that orbit a red dwarf star around 40 light-years from Earth. This suggests the telescope should be able to detect life or alien civilizations on exoplanets within 40 light-years of Earth. But the team believes JWST could possibly detect signs of extraterrestrial life up to 50 light-years from Earth.

Only around 20 exoplanets have been officially discovered within a 50-light-year radius of Earth, but based on the number of suspected stars in this region of space, experts predict that there may actually be as many as 4,000 exoplanets within JWST's reach, according to Project EDEN, an international astronomical collaboration dedicated to finding potentially habitable planets close to Earth.

However, this doesn't guarantee that JWST would be able to detect life on other planets.

Detecting biosignatures and technosignatures on other worlds "may prove challenging to interpret without contextual knowledge about the habitable environment," the researchers wrote. In this study, the team already knew which markers to look for, but on an exoplanet with different conditions and alternate potential life forms or technologies those life-signatures may not be as obvious, they added.

JWST has already made some interesting discoveries about exoplanets near Earth. The telescope spotted water on the Neptune-size exoplanet GJ 1214b, which is around 40 light-years from Earth, and found that TRAPPIST-1b, the second-closest exoplanet to the star in the TRAPPIST-1 system, likely has no atmosphere at all due to its extreme heat. The spacecraft also glimpsed a gigantic dust storm in the atmosphere of VHS 1256 b, a "super-Jupiter" exoplanet 40 light-years from Earth.

Closer to home, JWST has also detected giant geysers gushing out of Saturn's moon Enceladus, which could contain the chemical ingredients needed for life. And further out into the cosmos, the spacecraft has also glimpsed potentially life-giving carbon compounds in an infant star system more than 1,000 light-years from Earth.

Scientists have discovered and synthesized an entirely new isotope of the highly radioactive element uranium. But it might last only 40 minutes before decaying into other elements. 

The new isotope, uranium-241, has 92 protons (as all uranium isotopes do) and 149 neutrons, making it the first new neutron-rich isotope of uranium discovered since 1979. While atoms of a given element always have the same number of protons, different isotopes, or versions, of those elements may hold different numbers of neutrons in their nuclei. To be considered neutron-rich, an isotope must contain more neutrons than is common to that element.

Uranium is in the class of elements in the periodic table known as "actinides," which have proton counts between 89 and 103. All actinides are radioactive, but uranium is one of the four most radioactive elements, alongside radium, polonium and thorium.

"We measured the masses of 19 different actinide isotopes with a high precision of one part per million level, including the discovery and identification of the new uranium isotope," Toshitaka Niwase, a researcher at the High-energy Accelerator Research Organization (KEK) Wako Nuclear Science Center (WNSC) in Japan, told Live Science in an email. "This is the first new discovery of a uranium isotope on the neutron-rich side in over 40 years."

Related: Lightest known form of uranium created

Niwase is the lead author of a study on the new uranium isotope, which was published March 31 in the journal Physical Review Letters. 

Isotopes can be stable, meaning they keep their atomic configuration, or unstable, meaning they decay and break down into other elements by gaining or shedding protons. Decay rates are measured by an isotope's half-life, or the time it takes for half the material to decay into other elements. After two half-lives, a quarter of the material remains; after three, an eighth, and so on. 

The team hasn't yet measured the half-life of uranium-241, but theoretical estimates put it at around 40 minutes, Niwase said. This is somewhat short for a half-life. (For reference, the half-life of carbon-14 is 5,730 years, the half-life of the very unstable isotope technetium-99m is six hours and the half-life of francium-223 is 22 minutes. The fastest-decaying isotope, hydrogen-7, is half gone in just 10^-23 seconds.)

Niwase and colleagues created the uranium-241 by firing a sample of uranium-238 at platinum-198 nuclei at Japan's RIKEN accelerator. The two isotopes then swapped neutrons and protons — a phenomenon called "multinucleon transfer."

The team then measured the mass of the created isotopes by observing the time it took the resulting nuclei to travel a certain distance through a medium. The experiment also generated 18 new isotopes, all of which contained between 143 and 150 neutrons.

Niwase acknowledged that uranium-241 probably doesn't have many useful practical or scientific implementations, as the isotope is created in extremely small numbers.

Francium — the 87th element on the periodic table  — is a naturally occurring, but incredibly rare, radioactive element. It forms and decays extremely quickly, so it has no practical uses, and it is mostly used in scientific research.

However, the element has some intriguing properties: It is one of the only elements with no known stable form, and it has the largest atomic radius. 

Some elements on the periodic table are relatively abundant on Earth — hydrogen, helium, oxygen and carbon, for example — while others are far more elusive — promethium and thulium being the two most uncommon. Francium sits very much at the elusive end of the spectrum. 

"Francium-223 forms naturally during the radioactive decay of other elements, but it is estimated that there is only about 30 grams [1 ounce] of francium in the entire crust of the Earth at any one time," Christopher Barnett, a postdoctoral chemistry researcher at the University of Sydney, Australia, told Live Science in an email.

The element was first found in 1939 by French physicist Marguerite Perey, a prodigy of Marie Curie, who named the new discovery after Perey's homeland. This scarce element is considered to be one of the last naturally occurring elements discovered on Earth.

Francium fast facts

• Atomic number (number of protons in the nucleus): 87
• Atomic symbol (on the periodic table of elements): Fr
• Atomic weight (average mass of the atom): 223
• Density: unknown
• Phase at room temperature: solid
• Melting point: 80.6 degrees Fahrenheit (27 degrees Celsius)
• Boiling point: 1,250 F (677 C)

Uses and half-life

Because francium is so rare, scientists who want to study it first have to create it in nuclear reactions — either by “bombarding radium with neutrons, or thorium with protons," Barnett said.

Radium — atomic number 88 — is a radioactive, metallic chemical element, while thorium — atomic number 90 — is a naturally occurring radioactive metal. 

So, is francium useful? Have scientists found any practical applications for it? The simple answer is "no."

The most stable form of francium, according to Barnett, is francium-223, but even so "the half-life is so short that there are no known biological processes that use it," Barnett said. 

A half-life is a measure of how long it takes half of a sample of radioactive material to decay."One might expect the decay [of any element] to happen linearly and at a constant rate," Barnett said. "However, we have found that the rate of decay changes with the amount of substance left, such that the amount of material will halve in a constant time period."

Francium has a half-life of 22 minutes. So, if observing a sample of francium, half of the original amount will be left after 22 minutes. After another 22 minutes, there will be a quarter of the initial amount. It decays into radium-223 through beta decay, when an electron is emitted, or astatine-219 through alpha decay, when an "atom's nucleus sheds two protons and two neutrons in a packet that scientists call an alpha particle," according to JLab.

"Francium's half-life is extremely short for a radioactive element. For comparison, technetium-99m (which is used in medical imaging) has a half-life of six hours. Uranium-235 (the type used in nuclear reactors) has a half-life of 703,800,000 year[s]," Barnett said.

While francium's half-life is incredibly short and though it is very toxic due to its radioactivity, its characteristics make it a compelling subject to study.

"Francium has some interesting properties. It is used to develop and test theoretical models, and I can see these being very useful to test machine-learning-based physics and chemistry models." 

Barnett also noted that francium has never been observed with the naked eye because "there has never been enough isolated [francium] to know what it looks like." Barnett predicts that it will probably be "a silvery-grey metal," similar to other alkali metals, but admits "we just don't know for sure."

Is francium dangerous?

Given that francium can only be found in miniscule quantities, it doesn't pose much of a risk to humans, but what if, hypothetically speaking, someone was able to get hold of a large amount of it? Would francium be considered dangerous?

"There are two main risks with francium," Barnett said. "One is radioactivity. When it decays, it releases high energy particles that can ionise the surrounding tissues (causing burns), or break DNA strands (resulting in cancer). 

Yet, because francium is so hard to make, getting enough to cause a serious issue would be unlikely. The most ever isolated was less than 0.000000001% of that found in a brand-new smoke detector. Given its much shorter half-life, most of that radiation would be emitted in a few hours, compared to years for a smoke detector. 

"The other risk is that it is an alkali metal. It is — or rather would be, if we could get enough of it together — very reactive, and would likely catch fire most spectacularly."

Originally published on Live Science on Sept. 11, 2013, and rewritten on July 26, 2022.

Arsenic rose to infamy centuries ago as a nearly odorless, tasteless poison that was often used by and against the ruling classes in Europe during the Middle Ages and the Renaissance. 

But what is the history of arsenic poisoning, and how does it kill?

It turns out, an element that's vital to life also plays a role in making arsenic lethal.

What is the history of arsenic poisoning?

Arsenic is a naturally occurring element that is widely distributed in Earth's crust, according to the Centers for Disease Control and Prevention's (CDC) Agency for Toxic Substances and Disease Registry. Pure arsenic — which is a steel-gray, brittle solid — is typically found in the environment combined with other elements, such as oxygen, chlorine, sulfur, carbon and hydrogen, often resulting in white or colorless powders that have no smell or special taste. As such, you can't usually tell if arsenic is present in food, water or air.

Historically, arsenic was known as both the "king of poisons" and the "poison of kings," for its toxic power and its popularity among rulers who wanted to quietly do away with their rivals, according to a 2011 study published in the journal Toxicological Sciences. 

Stories abound describing arsenic's deadly use. For example, in biomedical historian James C. Whorton's book "The Arsenic Century" (Oxford University Press, 2010), Whorton recounted the legend of Roman emperor Nero ridding himself of his 13-year-old stepbrother and potential rival Britannicus by slipping arsenic into his soup.

Related: Why don't poisonous animals die from their own toxins?

Powerful and wealthy Italian families, such as the Medici and the Borgia, were also rumored to have used arsenic to eradicate their rivals, according to the Toxicological Sciences report. The use of arsenic in murder was common until the development in the 18th century of chemical methods of detecting arsenic poisoning, which involve looking for the element in hair, urine or nails, according to Britannica.

Nowadays, arsenic poisoning is more likely to be accidental than deliberate. People are most frequently exposed to arsenic through drinking water in areas where arsenic levels in dissolved minerals are naturally high, according to the CDC. Other sources of accidental arsenic exposure include contact with contaminated soil or dust, wood that has been preserved using arsenic compounds, or certain foods, such as rice and some fruit juices. (Rice absorbs an unusual amount of arsenic from the soil compared with other crops, according to the FDA; the agency notes that arsenic may make its way into apple and other juices due to naturally high levels of arsenic in soil and water, past use of arsenic-based pesticides in the United States and current use of such pesticides in other countries.)

What makes arsenic toxic?

Arsenic's toxicity stems from its proximity to phosphorus on the periodic table of elements. Because arsenic and phosphorus have similar atomic structures, they have similar properties. Both possess chemical keys that unlock access to cellular function. But whereas phosphorus is essential to life, arsenic is disruptive and deadly, Mark Jones, a chemistry consultant and fellow of the American Chemical Society, told Live Science.

Arsenic's similarity to phosphorus means that "arsenic can substitute very easily for phosphorus in many fundamental chemical reactions in biology and disrupt them," Jones said. "This means that arsenic can act like a broad-spectrum poison against insects, weeds and pretty much every life-form."

Related: What should you do if you're bitten by a venomous snake?

For example, phosphorus helps cells generate adenosine triphosphate (ATP), which is the main source of energy in all known organisms, according to the American Chemical Society. Arsenic can mimic phosphorus in chemical interactions where enzymes use oxygen to help liberate the energy stored in the sugar glucose and capture it within ATP. This can lead to arsenic disrupting the vital chemical reactions in which phosphorus takes part.

"You can think of enzymes and the chemicals they act upon as locks and keys," Jones said. "Arsenic is like a key that is not cut correctly — if it goes into a lock on a door, not only will it not unlock that door, it can get jammed in there and prevent another key from getting in to unlock that door. In this way, arsenic can block a lot of vital chemical pathways."

By chemically jamming cellular "locks," arsenic can harm nearly every organ in the human body. Large doses can lead to symptoms including vomiting, diarrhea, dehydration, shock, abnormal heart rhythms and multiple-organ failure, which may ultimately result in death, according to the CDC. Long-term exposure to high levels of arsenic in drinking water is linked to medical conditions such as skin disorders, an increased risk for diabetes, high blood pressure, and several types of cancer, including lung and skin cancers, the CDC says.

Individual susceptibility to arsenic poisoning varies widely; some people can tolerate doses of the element that would kill others, according to Britannica. In a 2018 study published in the journal Mammalian Genome, researchers reported that people's genes, diet and gut microbes may affect their chances for surviving an encounter with the deadly toxin.

Despite its deadly potential, arsenic poisoning is treatable if caught early, according to the Agency for Toxic Substances and Disease Registry. A key medicine is dimercaprol, which was developed by British scientists during World War II as an antidote to arsenic-based chemical weapons. The drug works by absorbing arsenic and neutralizing its toxicity, according to the National Library of Medicine. 

Although arsenic has a reputation for being deadly, it can also help cure disease, according to the Wellcome Library in England. In 1909, German chemist and Nobel Prize winner Paul Ehrlich and his colleagues developed an arsenic-loaded compound called Salvarsan, which became the first effective treatment for syphilis, according to the Science History Institute in Philadelphia. The principle behind how Salvarsan works, wherein a drug seeks out and destroys diseased cells, eventually found use in chemotherapy, Wellcome Library reported.

Originally published on Live Science.

One side of the moon is littered with far more craters than the other, and researchers finally know why: A massive asteroid that slammed into the moon around 4.3 billion years ago wreaked havoc in the moon's mantle, according to a new study. 

More than 9,000 visible craters pockmark the moon,  thanks to barrage of impacts from meteors, asteroids and comets over billions of years, according to the International Astronomical Union. However, these craters are not evenly distributed across the lunar surface. The far side of the moon, which people never see from Earth because the moon is tidally locked (meaning that it takes the same amount of time for the moon to rotate and orbit Earth), has a considerably higher concentration of craters than the visible nearside.

The nearside of the moon has fewer pits because the surface is covered in lunar maria — vast stretches of solid lava that we can see with the naked eye on Earth as dark patches on the moon. These lava fields likely covered up the craters that would otherwise have marked the moon's nearside. The far side of the moon has almost no lunar maria, which is why its craters are still visible.

Scientists have long suspected that lunar maria formed in the wake of a massive collision around 4.3 billion years ago. This collision created the South Pole–Aitken basin (SPA), a huge crater with a maximum width of around 1,600 miles (2,574 kilometers) and a maximum depth of 5.1 miles (8.2 km), which is the largest pit on the moon and the second largest confirmed impact crater in the solar system. However, until now researchers were unable to explain why only the nearside of the moon has lava fields. 

Related: How many space rocks hit the moon every year? 

The new study finds that the SPA impact created a unique phenomenon inside the moon's mantle, the layer of magma below the crust, that  affected only the nearside. 

"We know that big impacts like the one that formed SPA would create a lot of heat," lead author Matt Jones, a doctoral student of planetary science at Brown University, said in a statement. "The question is how that heat affects the moon's interior dynamics." 

Researchers already knew the nearside's lava fields originated within the moon's mantle, because lunar samples brought back by the Apollo missions contained radioactive, heat-generating elements such as potassium, phosphorus and thorium that are all suspected to be found in abundance within the lunar mantle, according to the statement.

In the new study, computer simulations revealed that the SPA impact would have created a heat plume within the mantle that pushed the radioactive elements toward the crust. The researchers repeated the simulation for a number of possible scenarios of the SPA impact, including direct hits and glancing blows, and found that regardless of how the asteroid hit, the mantle impacts would have only affected the nearside of the moon.

Put another way, when a space rock collided with the moon, it caused lava from the mantle to pour out on the nearside, burying many of its older impact craters. 

"What we show is that under any plausible conditions at the time that SPA formed, it ends up concentrating these heat-producing elements on the nearside," Jones said. "We expect that this contributed to the mantle melting that produced the lava flows we see on the surface."

The researchers are pleased to have solved what they described as "one of the most significant questions in lunar science," according to the statement.

"The SPA impact is one of the most significant events in lunar history," Jones said. Being able to better understand how it shaped the two sides of the moon we see today is "really exciting," he added.

The study was published online April 8 in the journal Science Advances.

Originally published on Live Science.

Editor's Note: This story was updated at 12:15 p.m. EDT to correct a description of tidal locking.

There are 118 elements on the periodic table, and you might think we are made up of many of them. But that's not the case; the complex systems that make up our bodies have a surprisingly simple elemental makeup. 

"Of 118 elements, the last thing I read was that 97% of our body's weight is just four elements," Steven Townsend, an organic chemist at Vanderbilt University, told Live Science. Those four elements are oxygen, carbon, hydrogen and nitrogen. So what, exactly, do these elements do? 

Oxygen definitively holds first place, making up 65% to 67% of the human body by weight, Townsend said. That's because our bodies are mostly water — around 50% to 60% — so most of the oxygen is the "O" in H2O. Oxygen is also critical for energy production and metabolism, or the chemical processes that occur within the body, according to a 2016 review in the journal Nature Reviews Cancer. 

Next up is carbon, which makes up around 18% to 19% of the body by weight, Townsend said, which isn't surprising considering carbon is a major component of most life on Earth. Carbon makes up the backbone of fats, carbohydrates and proteins, so this element is a major building block of the body and the foods we use to fuel it.

Related: How many organs are in the human body?

Coming in at No. 3 is hydrogen. There are more hydrogen atoms in the human body than any other element, but it makes up just 9% to 10% of our bodies by mass. Hydrogen is the other element in H2O, and there are two hydrogen atoms for every oxygen atom. Plus, hydrogen is a key component of proteins, carbohydrates and fats. But even though hydrogen holds the majority in number of atoms, it's significantly out-massed: Oxygen and carbon have atomic weights that are nearly 16 and 12 times that of hydrogen, respectively. In other words, it takes 16 hydrogen atoms to match the mass of one oxygen atom, so that's why hydrogen makes up only one-tenth of our mass.

Number four, at around 3%, is the most plentiful gas in Earth's atmosphere: nitrogen, according to Townsend. Nitrogen is critical to proteins and their building blocks, called amino acids. Nitrogen is also a major component of DNA and RNA, including of their nitrogenous bases. Both have cytosine, adenine and guanine; DNA has thymine, and RNA sports uracil. In other words, without nitrogen, the cells in your body couldn't store genetic information or replicate. A bonus element, the fifth most common in the human body, is calcium, which makes up 1% to 2% of humans by mass. More than 99% of that calcium is found in the bones and teeth, according to the National Institutes of Health.

However, elements other than these top five matter, too Townsend said. Take sodium, for instance. It makes up only 0.2% of the body, but "it's a super important mineral — it helps balance the fluids in the body," he said. If sodium is out of balance, people can have significant health problems, such as high blood pressure or loss of kidney function.

"[The human body is] intricate, but I think it's also amazing in its simplicity," Townsend said. "The complex interconnected systems that keep us alive — they all depend on a handful of elements. If you think about how complex the human body is, that's sort of magical."

Originally published on Live Science.

The element gold is a pirate's booty and an ingredient in microcircuits. It's been used to make jewelry since at least 4000 B.C. and to treat cancer only in recent decades. It's in the pot at the end of the rainbow and in the coating on astronaut visors. Gold is an element that bridges old and new — and myth and science — seamlessly.

Properties of gold

Gold, the 79th element on the Periodic Table of the Elements, is one of the more recognizable of the bunch. It is malleable and shiny, making it a good metalworking material. Chemically speaking, gold is a transition metal. Transition metals are unique, because they can bond with other elements using not just their outermost shell of electrons (the negatively charged particles that whirl around the nucleus of an atom), but also the outermost two shells. This happens because the large number of electrons in transition metals interferes with the usual orderly sorting of electrons into shells around the nucleus.

• Atomic Number (number of protons in the nucleus): 79
• Atomic Symbol (on the Periodic Table of Elements): Au
• Atomic Weight (average mass of the atom): 196.9665
• Density: 19.3 grams per cubic centimeter
• Phase at Room Temperature: Solid
• Melting Point: 1,947.7 degrees Fahrenheit (1,064.18 degrees C)
• Boiling Point: 5,162 degrees F (2,850 degrees C)
• Number of isotopes (atoms of the same element with a different number of neutrons): Between 18 and 59, depending on where the line for an isotope is drawn. Many artificially created gold isotopes are stable for microseconds or milliseconds before decaying into other elements. One stable isotope.
• Most common isotopes: Au-197, which makes up 100 percent of naturally occurring gold.

How is gold formed?

Gold represents a tiny fraction of the elements in the known universe. The reason for its rarity is owed to the incomprehensible amount of energy needed for its formation. Gold is formed in stars, but only in those that are exploding in giant supernovas, or incredibly dense ones that have come together in monstrously powerful collisions, according to the journal PNAS . 

Stars, such as our sun, generate energy through the power of fusion, where smaller elements are fused, or combined, together into heavier elements. To start with, a star may be mostly hydrogen, the smallest element. The process of fusion under immense pressure and heat in the star's core will generate helium. When hydrogen runs low and the star begins to reach the next phase of its life cycle, it will fuse helium into the next heavier element, and so on. 

This process continues until the element of iron, where the balance suddenly shifts. Because fusing iron does not create energy, it consumes it, according to the University of Oregon. With no means of generating internal energy to counteract its own immense pressure and gravity, the star begins to collapse onto itself. If the star is large enough the result is a supernova — a massive star explosion, according to NASA. Heavier elements are formed during the incredible energy generated during this process, including gold. 

Related: How can a star be older than the universe?

Gold throughout history

From Eastern Europe to the Middle East to the tombs of Egyptian Pharaohs, gold appears throughout the ancient world. Five thousand years ago, the massive Nile River was the key to the ancient Egyptian empire, according to the Australian government. Its water allowed a bounty of crops to be grown along its edge, keeping its citizens, and its armies, well fed. But there was also a shiny yellow metal that came running down the river, the element of gold. The Egyptians eagerly took this visually appealing treasure and found that because it was naturally pure and malleable, it required little refinement to be turned into mesmerizing decorations. 

Gold as a decoration didn't stop at ancient Egypt: A Stone Age woman found buried outside of London wore a strand of gold around her neck; Celts in the third century B.C. wore gold dental implants; a Chinese king who died in 128 B.C. was buried with gold-gilded chariots and thousands of other precious objects. 

Gold swiftly came to be a symbol, and unit, of wealth, and it has maintained this allure through time and around the globe. Several millennia after the Egyptian pharaohs and their tombs of gold, the Aztec Empire's gold riches were plundered by the Conquistadors who sought the valuable metal for their own. Later still, workers flocked to Western coast of the United States to take part in the California "gold rush", seeking their own fortunes, according to National Geographic. Therefore gold has driven humans to diplomacy, mass migrations, and even acts of genocide. Without this metal, our history would be quite different.

Related: Tenochtitlán: History of Aztec capital

Gold also plays a strong role in Australian history. In the late 19th century, so many flocked to the country to take part in its booming gold rush that the population of Australia tripled. Owing to its pervasive deposits, the country is still mined for the metal today, according to the Australian government. However, one company, named Evolution Mining, found a different treasure in their hunt for gold. When drilling into the Australian outback's surface in search of gold deposits, the miners instead unearthed sheets of stone that resembled "shatter cones," which form on the outer rims of impact craters. They followed this finding with advanced mapping techniques that allowed the team to confirm the uncovering of a 3.1-mile-wide (5 kilometers) meteorite crater, a finding even more rare than a lode of gold, according to Forbes.

What is a karat of gold?

Most gold jewelry isn’t made of pure gold. The amount of gold in a necklace or ring is measured on the karat scale. Pure gold is 24 karats. Bars of gold kept in Fort Knox and elsewhere around the world are considered to be 99.95 percent pure, 24-karat gold.

As metals are added to gold during jewelry-making, the gold becomes less fine and the number of karats drops. For example, 12 karat gold contains 50% gold and 50% alloys by weight.

The word karat comes from the carob seed. In ancient Asian bazaars, the seeds were used to balance scales that measured the weight of gold.

How much gold is in Fort Knox?

To keep up with the country's mounting gold reserves, the United States Bullion Depository opened at the Fort Knox U.S. Army Garrison in Kentucky in 1937. The first shipment of gold arrived from Philadelphia in trains surrounded by military troops.

Fort Knox is framed in steel with walls of concrete. Despite the defense of a 20-ton steel door, a dirty rumor in the 1970s suggested that the gold in Fort Knox was gone. To quell people's fears, the director of the United States Mint guided congress people and journalists through one room of the vault, and its 8-foot-tall stacks of 36,236 bars of gold.

The depository holds about half of U.S. Treasury's stored gold, according to the U.S. Mint. Each bar weighs 400 troy ounces (about 27.5 pounds), according to the U. S. Department of Treasury. One troy ounce equals about 1.1 avoirdupois ounces. The entire stockpile, as of 2021, weighs 147.3 million troy ounces, which is worth about $130 billion at today's prices. Fort Knox held a record amount of gold on Dec. 31, 1941, reaching a whopping 649.6 million ounces, the U.S. Mint reported.

Other important artifacts have also "seen" the insides of Fort Knox. For instance, during WWII, the Declaration of Independence, Constitution and Bill of Rights were sealed inside for protection, being returned in 1944 to Washington, D.C. Other items stored there at some point in history, according to the U.S. Mint include: the Magna Carta; the crown, sword, scepter, orb and cape of St. Stephen, the King of Hungary.

What is fool's gold?

Pyrite, the inferior mineral nicknamed fool's gold, only mimics gold in looks. Pyrite is more common, harder, and more brittle than gold. When crushed into powder, it looks greenish-black, whereas real gold powder is yellow. Pyrite contains sulfur and iron. During World War II, it was mined to produce sulfuric acid, an industrial chemical. Today, it is used in car batteries, appliances, jewelry and machinery.

Although fool's gold can be a disappointing find, it is often discovered near sources of copper and real gold. So perhaps, miner who stops digging once they have a piece of pyrite in hand is the real fool.

Fun facts about gold

• Two-thirds of the world's gold is mined in South Africa, according to Lawrence Livermore National Laboratory.
• Seventy-eight percent of the world's yearly supply of gold is used in jewelry, according to the AMNH. The rest goes to electronics and dental and medical uses.
• The atomic symbol of gold, Au, comes from the Latin word for gold, aurum.
• Astronaut helmets were equipped with a visor coated with a thin layer of gold. The gold blocks harmful ultraviolet rays from the sun.
• The world's largest gold crystal is the size of a golf ball and comes from Venezuela. The 7.7-ounce (217.78 grams) crystal is worth about $1.5 million.
• Earthquakes can create gold: A 2013 study in the journal Nature Geoscience found that during earthquakes, water in faults and fractures vaporizes, leaving gold behind.
• The first purely gold coins were manufactured in the Asia Minor kingdom of Lydia in 560 B.C., according to the National Mining Association.
• Gold has a number of artificial, unstable isotopes (the exact number depends on the scientist you consult), but occurs naturally only as Au-197.
• You can eat gold … if you really want to. Gourmet shops sell edible gold leaf and flakes that add glitter to everything from pastries to vodka to olive oil. Don't fear for your stomach: The gold isn't digested and just passes right through, according to Edible Gold, a company that sells gold leaf.

How is gold used?

Gold is also used in medicine. The radioactive gold isotope Au-198 can be injected directly into the site of a tumor, where its radiation can destroy tumor cells without much spillover to the rest of the body. In 2012, researchers reported in the journal Proceedings of the National Academy of Sciences that they could link nanoparticles of Au-198 with a compound found in tea leaves to treat prostate cancer. The tea compound is attracted to the tumor cells, keeping the nanoparticles glued to the right spot for several weeks while the radiation treatment occurs. (The method has yet to be tested on humans.)

In some cases, gold nanoparticles are the only way a drug can work. The anti-cancer drug TNF-alpha kills cancer very effectively. Unfortunately, it's also incredibly toxic to healthy cells. However, clinical trials now underway have found that linking TNF-alpha drugs to gold nanoparticles can successfully treat tumors, because the drugs hit their targets directly, according to Benchmarks, an online publication of the National Cancer Institute.

There's just one problem with humanity's continued love affair with gold: Getting it out of the ground. About 83% of the 2,700 tons of gold mined each year is extracted using a process called gold cyanidation, said Zhichang Liu, a postdoctoral researcher in chemistry at Northwestern University in Illinois. This process uses cyanide to leach gold out of the rock that holds it. Unfortunately, cyanide is toxic, and the process is anything but environmentally friendly.

There could be hope for lovers of gold baubles (and electronic circuits and nanomedicine), however. In 2013 Liu and his colleagues reported in the journal Nature Communications that they'd stumbled upon a way to extract gold from ore with benign starch rather than toxic cyanide.

"Actually, we found this method by accident," Liu told Live Science. While trying to fabricate a porous material, the researchers mixed a starch called alpha-Cyclodextrin with gold salts (charged molecules of gold). To their surprise, the gold precipitated out of the solution rapidly.

The team has since patented the method, which easily extracts gold at more than 97% purity in one step, Liu said. They're now working with investors to scale up the process. "Hopefully, we can find a nice, green way to replace the cyanidation process," Liu said.

Additional resources

• Learn more about the California Gold Rush on The History Channel.
• Michael D. Bordo, a professor of economics at Rutgers University, provides a thorough look at the Gold Standard.
• DK and Smithsonian put together a kids book with loads of visualizations about the Periodic Table.
• Learn more about gold mining around the world on the World Gold Council page.

Time for a science refresher course? One in 5 Americans can't name a single element on the Periodic Table.

Most Americans surveyed (59 percent) couldn't name more than 10 elements of the 118 that grace the Periodic Table. This may have been a result of the way the question was asked: Americans probably do know many elements' names (gold is one; so are other household names like silver, tin, lead, oxygen, helium and calcium), but may not realize that they are, in fact, elements.

The new survey comes courtesy the Philadelphia-based nonprofit the Science History Institute. It was administered via the consulting organization YouGov, which polled 1,263 adults online and weighted the responses to be representative of the demographics of U.S. adults. [6 Important Elements You've Never Heard Of]

Periodic table primer

Elements are the basic building blocks of matter; substances earn spots on the Periodic Table because they can't be broken down into anything simpler. That chart that graces science classrooms worldwide dates back to 1869, when Russian chemist Dmitri Mendeleev presented his brand-new way of organizing the known elements by atomic mass (the number of protons and neutrons in an atom) and valence (the maximum number of electrons in an atom's outer shell, which are available for bonding with other atoms).

The Periodic Table was most recently updated in 2016, when four new elements made their debut. For those who'd like to ace the next Science History Institute survey, their names are nihonium, moscovium, tennessine and oganesson. These elements are superheavy, with 113, 115, 117 and 1118 protons in their nuclei, respectively. That means they're very unstable. They don't occur naturally, and when they are created in the lab, they rapidly decay into other, more stable elements.

The survey found that 57 percent of Americans believe in the importance of science, and 45 percent think it is important for them to stay up-to-date on scientific developments. But there were gaps in basic scientific knowledge. Seventeen percent of Americans said they feel it is intimidating to stay up-to-date, and 24 percent said they wish science information was more accessible.

Rare-earth elements

The survey also drives home a dearth of understanding of rare-earth elements. Twenty-six percent of those surveyed had not heard this term, and 35 percent had heard it, but had no idea what it meant.

The rare-earth elements are 17 elements with atomic numbers 57 through 71, plus 21 and 39. They're metals with similar properties, and they're important components of a lot of modern tech, from portable electronics to fuel cells to lasers. They get their name because they're rarely found in concentrated deposits, but they are actually quite common globally. Their names? Scandium, yttrium, lanthanum, cerium, praseodymium, neodymium, promethium, samarium, europium, gadolinium, terbium, dysprosium, holmium, erbium, thulium, ytterbium and lutetium (say that five times fast).

The survey found that people are very interested in the technology made possible by these elements. Fifty-four percent said they couldn't live without the internet, and 41 percent said they couldn't live without their smartphones. About 1 in 3 said clean-energy tech and climate-change-battling advances will be the most important technology of the future, while 20 percent voted for health technology and 18 percent for communication technology as making the biggest future impact. All rely on rare-earth elements.

• Elementary, My Dear: 8 Little-Known Elements
• What's That? Your Physics Questions Answered
• The 18 Biggest Unsolved Mysteries in Physics

Originally published on Live Science.

Two astronomers think they've pinpointed the ancient stellar collision that gave our solar system its cache of precious gold and platinum — some of it, anyway.

In a new study published May 1 in the journal Nature, the duo analyzed the remnants of radioactive isotopes, or versions of molecules with different numbers of neutrons, in a very old meteorite. Then, they compared those values with isotope ratios produced by a computer simulation of neutron star mergers — cataclysmic stellar collisions that can cause ripples in the fabric of space-time. [15 Unforgettable Images of Stars]

The researchers found that a single neutron star collision, starting about 100 million years before our solar system formed and located 1,000 light-years away, may have provided our cosmic neighborhood many of the elements heavier than iron, which has 26 protons. This includes about 70% of our early solar system's curium atoms and 40% of its plutonium atoms, plus many millions of pounds of precious metals like gold and platinum. In total, this single ancient star crash may have given our solar system 0.3% of all its heavy elements, the researchers found — and we carry some of them around with us every day.

He added that, if you wear a gold or platinum wedding ring, you're also wearing a bit of the explosive cosmic past. "About 10 milligrams [0.00035 ounces] of it likely formed 4.6 billion years ago," Bartos said.

There's gold in them thar stars

How does a star make a wedding ring? It takes an epic cosmic explosion (and a few billion years of patience).

Elements like plutonium, gold, platinum and others heavier than iron are created in a process called rapid neutron capture (also called the r-process), in which an atomic nucleus quickly gloms on to a bunch of free neutrons before the nucleus has time to radioactively decay. This process occurs only as a result of the universe's most extreme events — in stellar explosions called supernovas or colliding neutron stars — but scientists disagree about which of those two phenomena is chiefly responsible for the production of heavy elements in the universe.

In their new study, Bartos and his colleague Szabolcs Marka (of Columbia University in New York) make an argument for neutron stars being the predominant source of heavy elements in the solar system. To do so, they compared the radioactive elements preserved in an ancient meteorite with numerical simulations of neutron star mergers at various points in space-time around the Milky Way.

"The meteor contained the remnant of radioactive isotopes produced by neutron star mergers," Bartos told Live Science in an email. "While they decayed a long time ago, they could be used to reconstruct the amount of the original radioactive isotope at the time when the solar system was formed."

The meteorite in question contained decayed isotopes of plutonium , uranium and curium atoms, which the authors of a 2016 study in the journal Science Advances used to estimate the amounts of these elements present in the early solar system. Bartos and Marka plugged those values into a computer model to figure out how many neutron star mergers it would take to fill the solar system with the correct amounts of those elements.

A casual cataclysm

It turns out that a single neutron star merger would do the trick, if it happened close enough to our solar system — within 1,000 light-years, or about 1% of the diameter of the Milky Way.

Neutron star mergers are thought to be pretty rare in our galaxy, occurring only a few times every million years, the researchers wrote. Supernovas, on the other hand, are much more common; according to a 2006 study from the European Space Agency, a massive star explodes in our galaxy once every 50 years or so.

That supernova rate is much too high to account for the levels of heavy elements observed in early solar system meteors, Bartos and Marka concluded, ruling them out as the likely source of those elements. A single nearby neutron star merger, however, fits the story perfectly.

According to Bartos, these results "shed bright light" on the explosive events that helped make our solar system what it is.

• 6 Cosmic Catastrophes That Could Wipe Out Earth
• The 12 Strangest Objects in the Universe
• 9 Strange Excuses for Why We Haven't Found Aliens Yet

Originally published on Live Science.

A new kind of matter can be both solid and liquid at once.

In this chain-melted state, molten and solid layers intertwine at the atomic level. Recently, using computer simulations, researchers coaxed virtual potassium into a chain-melted state by exposing the metal to conditions of extreme temperature and pressure, the scientists reported in a new study.

What's more, this dual state persisted even through dramatic changes in the experiments' conditions within the simulation. This evidence also showed that the chain-melted state is a stable type of matter and not merely a transition between solid and liquid. [The 18 Biggest Unsolved Mysteries in Physics]

These experiments were conducted at the atomic level in a virtual environment, but what might it be like to hold an object in this peculiar state?

"It would look and feel like a solid, so you could pick it up, then there's a liquid part in it that could leak out," study co-author Andreas Hermann, a reader in computational physics with the University of Edinburgh's School of Physics and Astronomy in Scotland, told Live Science.

"But once the liquid gets lost from the material, some of the solid part would melt to replenish it," Hermann said.

The researchers had already demonstrated in a prior study that potassium, a highly reactive metal, was a little weird. They showed that under high pressure, potassium forms an unusual crystal structure of two different, interwoven lattices, "going from a very simple atomic arrangement to something very complicated," Hermann said.

For the new study, the scientists ran simulations that subjected potassium to high temperatures in addition to high pressure. Incorporating machine learning into the simulations greatly increased the number of atoms — 20,000 at once in this case — that the study authors could test.

In the new simulations, when things heated up, potassium did something very strange. After its atoms formed an interlocked lattice structure, the atoms in one lattice were strongly connected, maintaining a solid state. But the signal from the other lattice vanished, indicating disorder in the atoms, the study authors noted.

In other words, these atoms became liquid while their immediate atomic neighbors remained solid, creating a state that is neither truly solid nor liquid, but a mixture of both, "interconnected on the atomic level," Hermann said.

Once the potassium samples reached this dual state, they lingered as part-liquid and part-solid even after the heat was turned up hundreds of degrees, according to Hermann.

Other studies have shown that potassium isn't the only element that develops two intertwined lattices of atoms under intense pressure, and these elements — "neighbors of potassium and elsewhere on the periodic table" — may also be capable of attaining a part-liquid and part-solid state, Hermann said.

And the machine learning system that the study authors developed to examine potassium could also be used with other substances, to decode how extreme conditions affect them at the atomic level.

"This is the proof of principle: a computationally cheap technique that can describe materials across a wide range of pressures and temperatures, including some very exotic states like the one we wrote this paper about," Hermann said. "That's our aim, to move on to other materials where we can answer different materials-science related questions."

The findings will be published online in a forthcoming issue of the journal Proceedings of the National Academies of Science.

• The Mysterious Physics of 7 Everyday Things
• Image: Inside the World's Top Physics Labs
• Twisted Physics: 7 Mind-Blowing Findings

Originally published on Live Science.

Our sun is a lifeless, fiery ball of gas fueled by a nuclear inferno. Earth, meanwhile, is a rocky, layered planet covered by water and teeming with life. Nevertheless, the elemental composition of these two celestial bodies is surprisingly similar.

The elements in the sun and Earth are pretty much the same, though Earth had less of the sun's more volatile elements, which evaporate at high temperatures, a new analysis reveals.

This suggests that Earth formed from material in the solar nebula — the cloud of dust and gas that shaped the sun — but volatile elements such as helium, hydrogen, oxygen and nitrogen were stripped away during our planet's formation. The tools used in the current study could also help reveal the composition of exoplanets orbiting distant stars, the study authors reported. [Spaced Out! 101 Astronomy Images That Will Blow Your Mind]

First, the researchers analyzed elements that appeared in rocky meteorites that fell to Earth, known as chondrites. Chondrites, which also formed in the protosolar nebula, are often used as proxies for understanding the sun's chemical makeup, the researchers wrote.

They also evaluated the sun's elemental composition from observations of radiation in the sun's photosphere — the outer "shell" that emanates light — and incorporated data from solar turbulence and theoretical models.

Though the most abundant elements in the sun are hydrogen and helium, the researchers discovered a total of 60 elements was abundant in both meteorites and photosphere; these elements were probably also plentiful in the protosolar nebula before the sun's birth, according to the study.

Then, the scientists compared their results to the elemental composition of Earth's core and primitive mantle, which can be gleaned through a combination of mathematical models, seismic data and rock samples. They found that while Earth shared most of the same elements as chondrites and the sun, Earth had "devolatilized" — lost volatile elements over time — and that this was "an inherent process" as the inner solar system took shape, the researchers wrote.

"This comparison yields a wealth of information about the way the Earth formed," study co-author Trevor Ireland, a professor of geochemistry and cosmochemistry with the Research School of Earth Sciences at the Australian National University (ANU) in Canberra, said in a statement.

Similar evaluations could be done for planets orbiting stars other than our sun.

"Rocky exoplanets are almost certainly devolatilized pieces of the stellar nebulae out of which they and their host stars formed," the researchers wrote in the study.

Pinpointing the elemental makeup of faraway exoplanets will play an important part in determining if they can support human life, lead study author Haiyang Wang, a doctoral candidate with ANU's Research School of Astronomy and Astrophysics, said in the statement.

"The composition of a rocky planet is one of the most important missing pieces in our efforts to find out whether a planet is habitable or not," Wang said.

The findings appeared online March 14 in the preprint journal arXiv, and will be published in a forthcoming issue of the journal Icarus.

• Shine On: Photos of Dazzling Mineral Specimens
• What Will Happen to Earth When the Sun Dies?
• Top 10 Questions About Earth

Originally published on Live Science.

Our world is full of chemicals that shouldn't exist.

Lighter elements, like carbon and oxygen and helium, exist because of intense fusion energies crushing protons together inside stars. But elements from cobalt to nickel to copper, up through iodine and xenon, and including uranium and plutonium, are just too heavy to be produced by stellar fusion. Even the core of the biggest, brightest sun isn't hot and pressurized enough to make anything heavier than iron.

And yet, those chemicals are abundant in the universe. Something is making them. [Elementary, My Dear: 8 Elements You Never Heard Of]

The classic story was that supernovae — the explosions that tear some stars apart at the end of their lives — are the culprit. Those explosions should briefly reach energies intense enough to create the heavier elements. The dominant theory for how this happens is turbulence. As the supernova tosses material into the universe, the theory goes, ripples of turbulence pass through its winds, briefly compressing outflung stellar material with enough force to slam even fusion-resistant iron atoms into other atoms and form heavier elements.

But a new fluid dynamics model suggests that this is all wrong.

"In order to initiate this process we need to have some sort of excess of energy," said study lead author Snezhana Abarzhi, a materials scientist at the University of Western Australia in Perth. "People have believed for many years that this sort of excess might be created by violent, fast processes, which might essentially be turbulent processes," she told Live Science.

But Abarzhi and her co-authors developed a model of the fluids in a supernova that suggest something else — something smaller — might be going on. They presented their findings earlier this month in Boston, at the American Physical Society March meeting, and also published their findings Nov. 26, 2018 in the journal Proceedings of the National Academy of Sciences.

In a supernova, stellar material blasts away from the star’s core at high speed. But all that material is flowing outward at about the same speed. So relative to one another, the molecules in this stream of stellar material aren't moving all that fast. While there may be the occasional ripple or eddy, there's not enough turbulence to create molecules past iron on the periodic table.

Instead, Abarzhi and her team found that fusion likely takes place in isolated hotspots within the supernova.

When a star explodes, she explained, the explosion isn't perfectly symmetrical. The star itself has density irregularities in the moment before an explosion, and the forces blasting it apart are also a bit irregular.

Those irregularities produce ultradense, ultrahot regions within the already-hot fluid of the exploding star. Instead of violent ripples shaking the whole mass, the supernova’s pressures and energies get especially concentrated in small parts of the exploding mass. These regions become brief chemical factories more powerful than anything that exists in a typical star.

And that, Abarzhi and her team suggest, is where all the heavy elements in the universe come from.

The big caveat here is that this is a single result and a single paper. To get there, the researchers relied on pen-and-paper work, as well as computer models, Abarzhi said. To confirm or refute these results, astronomers will have to match them against the actual chemical signatures of supernovae in the universe -- gas clouds and other remainders of a stellar explosion.

But it seems like scientists are a bit closer to understanding how much of the material all around us, including inside our own bodies, gets made.

• Gallery: Our Amazing Sun
• Fiery Folklore: 5 Dazzling Sun Myths
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Originally published on Live Science.

On Feb. 17, 1869, Russian chemist Dmitri Mendeleev published his first attempt to sort the building blocks of life into orderly groups. Now, 150 years later, we know the fruits of his labor as the Periodic Table of Elements — a quintessential piece of classroom wall art and indispensable research tool to anyone who's ever picked up a beaker.

As you can see for yourself in the hand-scrawled draft above, Mendeleev's first table looked very different than the one we know today. In 1869, only 63 elements were known (compared with the 118 elements we have identified today). As a student at Heidelberg University in Germany and later as a professor at St. Petersburg University, Mendeleev realized that by grouping elements according to their atomic weights, certain types of elements periodically occurred. [Elementary, My Dear: 8 Little-Known Elements]

Mendeleev honed this "periodic system," as he called it, by writing down the names, masses and properties of each known element on a set of cards. According to science historian Mike Sutton of Chemistry World, Mendeleev then laid these cards down before him — solitaire-like — and started shuffling them around until he found an order that made sense.

Ultimately, Mendeleev's eureka moment came to him in a dream, Sutton wrote. When he awoke, he arranged his element cards in vertical columns in order of increasing atomic weight, starting a fresh column to group elements with similar properties into the same horizontal row. With these guiding principles, he eventually created the world's first Periodic Table.

Mendeleev was so confident in his system that he left gaps for undiscovered elements, and even predicted (correctly) the properties of three of those elements. Those three elements — known now as gallium, scandium and germanium — were discovered within the next three years and matched Mendeleev's predictions, helping to solidify the reputation of his table, Sutton reported.

The table wasn't perfect (Mendeleev was unable to locate hydrogen using his system, for example), but it laid a solid groundwork for generations of chemists to build upon over the next 150 years.

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Originally published on Live Science.

The periodic table of elements is a familiar sight to anyone who's ever sat in a chemistry classroom — and apparently, that's been the case for nearly 150 years.

Conservators at the University of St. Andrews in Scotland announced the discovery of what they say is the world's oldest surviving example of a classroom periodic table of elements, dating to 1885.

The old classroom poster, printed in German on linen-backed paper, was discovered in 2014 while staffers from the university's School of Chemistry were cleaning out their storage room, according to a news release from the university. Among the clutter of decades-old lab equipment and chemical vials, staffers found an old cache of huge, rolled-up teaching charts.

One of the scrolls contained the aforementioned periodic table — inked onto paper so old that it started to crumble at the touch.

School records showed that the chart was purchased in Vienna by a St. Andrews chemistry professor in 1888, and the table likely hung in his classroom until his 1909 retirement. Researchers were able to further narrow down the poster's print date by looking at the elements represented (and those left out) on the chart. For example, "both gallium and scandium, discovered in 1875 and 1879, respectively, are present, while germanium, discovered in 1886, is not," the news release said.

According to the university, this table seems to be the only one from its period surviving anywhere in Europe.

In any case, the old chart dates closely to the conception of the periodic table, itself. Russian chemist Dmitri Mendeleev developed the world's first periodic table when, following days of work and a vivid dream, he ordered the known elements according to their atomic mass and ability to bond with other elements. Mendeleev presented his findings to the Russian Chemical Society in 1869, and the first periodic tables were published soon after.

According to professor David O'Hagan, former head of chemistry at the University of St Andrews, the "remarkable" table will be available for research at the university and will go on public display later this year.

"We have a number of events planned in 2019, which has been designated [the] international year of the periodic table by the United Nations, to coincide with the 150th anniversary of the table's creation by Dmitri Mendeleev," O'Hagan said in the statement.

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Originally published on Live Science.

Polonium (Po) is a very rare and highly volatile radioactive metal. Before Polish-French physicist Marie Curie's discovery of polonium in 1898, uranium and thorium were the only known radioactive elements. Curie named polonium after her homeland, Poland.

Polonium is of little use to humans, with the exception of some menacing applications: It was used as a trigger in the first atomic bomb and is also a suspected poison in a couple of high-profile deaths.

In commercial applications, polonium is occasionally used to remove static electricity in machinery or dust from photographic film. It can also be used as a lightweight heat source for thermoelectric power in space satellites.

Classification

Polonium is located in Group 16 and period 6 in the periodic table of elements. It's classified as a metal because polonium's electrical conductivity decreases as its temperature rises, according to the Royal Society of Chemistry.

The element is the heaviest metal of the chalcogens, a group of elements also known as the "oxygen family." All chalcogens are found in copper ores. Other elements in the chalcogen group include oxygen, sulfur, selenium and tellurium.

There are 33 known isotopes (atoms of the same element with a different number of neutrons) of polonium, and all are radioactive. This element's radioactive instability is what makes it a fitting candidate for use in atomic bombs.

Physical characteristics

• Atomic number (number of protons in the nucleus): 84
• Atomic symbol (on the periodic table of the elements): Po
• Atomic weight (average mass of the atom): 209
• Density: 9.32 grams per cubic centimeter
• Phase at room temperature: Solid
• Melting point: 489.2 degrees Fahrenheit (254 degrees Celsius)
• Boiling point: 1,763.6 degrees F (962 degrees C)
• Most common isotope: Po-210 which has a half-life of only 138 days

Discovery

When Curie and her husband, Pierre Curie, discovered polonium, they were searching for the source of radioactivity in a naturally occurring, uranium-rich ore called pitchblende.

The two noticed that the unrefined pitchblende was more radioactive than the uranium that had been separated from it. So, they reasoned that the pitchblende must be harboring at least one other radioactive element.

The Curies purchased loads of pitchblende so that they could chemically separate the compounds in the minerals. After months of painstaking work, they finally isolated the radioactive element: a substance 400 times more radioactive than uranium, according to the International Union of Pure and Applied Chemistry (IUPAC).

Extracting polonium was challenging because there was such a miniscule amount; 1 ton of uranium ore contains only about 100 micrograms (0.0001 grams) of polonium.

Nonetheless, the Curies were able to pull out the isotope we now know as polonium-209, according to the Royal Society of Chemistry.

Sources

Traces of Po-210 can be found in the soil and air. For example, Po-210 is produced during the decay of radon-222 gas, which a result of decaying radium. In turn, radium is a decay product of uranium, which is present in almost all rocks and soil formed from rocks.

Lichens are able to absorb polonium directly from the atmosphere. In northern areas, people who eat reindeer can have higher concentrations of polonium in their blood, because reindeer eat lichens, according to Smithsonian.com.

Polonium is considered a rare natural element. Although it is found in uranium ores, it's not economical to extract as there are only around 100 micrograms of polonium in 1 ton (0.9 metric tons) of uranium ore, according to the Jefferson Lab.

Instead, polonium is obtained by bombarding bismuth-209 (a stable isotope) with neutrons in a nuclear reactor. This creates radioactive bismuth-210, which then decays into polonium through a process called beta decay, according to the Royal Society of Chemistry.

The United States Nuclear Regulatory Commission estimates that only around 100 grams (3.5 ounces) of polonium-210 is produced worldwide each year.

Commercial uses

Because of its high radioactivity, polonium has few commercial applications. Among the element's limited uses are eliminating static electricity in machinery and removing dust from photographic film. In both applications, the polonium must be carefully sealed to protect the user.

The element is also used as a lightweight heat source for thermoelectric power in satellites and other spacecraft. That's because polonium decays rapidly, and as it does, it releases a large amount of energy in the form of heat. Just a single gram of polonium will reach a temperature of 500 degrees C (932 degrees F) as it degrades, according to the Royal Society of Chemistry.

Atomic bomb

During the middle of World War II, the Army Corps of Engineers began to organize the Manhattan Engineer District, a top-secret research and development program that would ultimately produce the world's first nuclear weapons.

Before the 1940s, there was no reason to isolate polonium in its pure form or to produce it in any substantial quantity, because there was no known use for it and very little was known about it. But the district's engineers began studying polonium and found the element to be an important ingredient for their nuclear weapon.

A combination of polonium and beryllium, another rare element, acted as the bomb's initiator, according to the Atomic Heritage Foundation.

After the war, the polonium research project was transferred to Mound Laboratory in Miamisburg, Ohio. Completed in 1949, Mound Lab was the first permanent Atomic Energy Commission facility for nuclear weapons development.

Poisoning

Polonium is toxic to humans, even in very small amounts.

The first person to die of polonium poisoning may have been Marie Curie's daughter Irène Joliot-Curie. In 1946, a polonium capsule exploded on her lab bench which may have been the reason she contracted leukemia and died 10 years later, according to Smithsonian.com.

Polonium poisoning was also what killed Alexander Litvinenko, a former Russian spy who had been living in London in 2006 after claiming political asylum.

Poisoning was also suspected in the 2004 death of the Palestinian leader Yasser Arafat, as surprisingly high levels of polonium-210 were detected on his clothes, according to The Wall Street Journal.

A 2011 study published in the journal Nicotine & Tobacco Research found that tobacco companies have been aware that cigarettes and other tobacco-containing products contain low levels of polonium. The study's authors calculated that the radioactivity from polonium in cigarettes is responsible for up to 138 deaths for every 1,000 smokers over a period of 25 years.

Other research has shown that twice as much polonium is found in the ribs of smokers as in those of nonsmokers, according to the U.S. National Institute of Health's Toxicology Data Network.

Further reading:

• Frequently Asked Questions about Polonium 210, from the CDC.
• Six Secrets of Polonium, from Smithsonian magazine.
• NIH's Hazardous Substances Data Bank entry for radioactive polonium.

The lightest known metal can also lighten your mood. Lithium, atomic number 3, is an element of many uses. It's used in the manufacture of aircraft and in certain batteries. It's also used in mental health: Lithium carbonate is a common treatment of bipolar disorder, helping to stabilize wild mood swings caused by the illness. 

Lithium has a flashy discovery story — literally. A Brazilian naturalist and statesman, Jozé Bonifácio de Andralda e Silva, discovered the mineral petalite (LiAISi4O10) on the Swedish isle Utö in the 1790s, according to the Royal Society of Chemistry (RSC). The mineral is white to gray, but when thrown into fire, it flares bright crimson. 

In 1817, Swedish chemist Johan August Arfwedson discovered that petalite contained a previously unknown element. He wasn't able to isolate the metal entirely, but he did isolate one of its salts. The name, lithium, comes from "lithos," the Greek word for "stone." 

It took until 1855 for someone to isolate lithium: British chemist Augustus Matthiessen and German chemist Robert Bunsen ran a current through lithium chloride to separate the element. 

Physical properties

According to the Jefferson National Linear Accelerator Laboratory, the properties of lithium are:

• Atomic number (number of protons in the nucleus): 3
• Atomic symbol (on the Periodic Table of Elements): Li
• Atomic weight (average mass of the atom): 6.941
• Density: 0.534 grams per cubic centimeter
• Phase at room temperature: Solid
• Melting point: 356.9 degrees Fahrenheit (180.5 degrees Celsius)
• Boiling point:  2448 degrees Fahrenheit (1342 degrees Celsius)
• Number of isotopes (atoms of the same element with a different number of neutrons): 10; 2 stable
• Most common isotopes: Li-7 (92.41 percent natural abundance), Li-6 (7.59 percent natural abundance)

The brain on lithium

Lithium is a special metal in many ways. It's light and soft — so soft that it can be cut with a kitchen knife and so low in density that it floats on water. It's also solid at a wide range of temperatures, with one of the lowest melting points of all metals and a high boiling point. 

Like its fellow alkali metal, sodium, lithium reacts with water in showy form. The combo of Li and H2O forms lithium hydroxide and hydrogen, which typically bursts into red flame. 

Lithium makes up a mere 0.0007 percent of the Earth's crust, according to the Jefferson Lab, and it's only found locked up in minerals and salts. Those salts have the power to change the brain: Lithium salts were the first drugs approved by the Food and Drug Administration to treat mania and depression, according to the National Institute of Mental Health.

Today, lithium carbonate is the compound most often sold as a pharmaceutical. No one knows exactly how lithium works to stabilize mood. Studies show multiple effects on the nervous system. In 2008, for example, researchers reported in the journal Cell that lithium interrupts the activity of a receptor for the neurotransmitter dopamine. It also appears to plump up brain volume, according to a 2011 study in the journal Biological Psychiatry (though this research is hotly contested).

In a study with worms, biologists at MIT found that lithium inhibits a key protein in the worms' brain, making neurons linked to an avoidance behavior go dormant. Essentially, the worms stopped avoiding harmful bacteria without that protein. The findings, which would need to be replicated in humans, suggest the element silences certain neurons in the brain and may have a calming effect, the researchers reported in 2016 in the journal Current Biology.

Lithium in space

Lithium, as well as the first and second lightest chemical elements (hydrogen and helium, respectively), are the only elements created at the birth of the universe, according to NASA. However, according to the Big Bang Theory, the universe should hold three times as much lithium as can be accounted for in the oldest stars, an issue called the missing lithium problem. This "missing lithium" discovery was first made in the 1980s, said Pasquale Serpico, a cosmologist at the National Center for Scientific Research (CNRS) and the University of Savoy Mont Blanc in France. It created a "tension," Serpico said, between what the Big Bang data and the observations of stars were telling researchers about lithium's abundance. 

Astrophysicists continue to conduct research to find this "missing" lithium or to explain why it's missing. In fact, researchers recently found a giant star holding 3,000 times more lithium than normal "giants," they reported in August 2018 in the journal Nature Astronomy. They came up with two possible explanations: the giant star swallowed its planet, absorbing the onboard lithium; the lithium also may have formed inside the star, reaching its surface before the heat of the deep layers vaporized it, according to a statement on the finding.

More about lithium

• Lithium-ion batteries are the key to lightweight, rechargeable power for laptops, phones and other digital devices. According to the U.S. Geological Survey, Argentina and Chile increased their lithium production 15 percent each in 2014 alone to meet the growing demand. Worldwide, production jumped 6 percent that year. 
• Lithium and another battery component, cobalt, could become scarce as demand increases, Stefano Passerini and Daniel Buchholz, both at the Helmholtz Institute Ulm in Germany, said in a statement describing their analysis of the future availability of those elements published in 2018 in the journal Nature Reviews Materials. In addition, both are concentrated in less politically stable countries, the study revealed. As such, the researchers urged the development of new battery technologies based on other, non-toxic elements.
• The United States has one lithium mine, in Nevada, according to the USGS. Chile and Australia produce the most lithium in the world.
• Naturally occurring lithium in drinking water correlates with lower levels of suicide, according to a 2009 study that highlights lithium's role in the brain. But psychiatrists are careful about prescribing lithium in high doses, particularly because it can pass through the placenta and have unknown effects on the developing fetus.  
• On a lighter note, the element is part of celebratory fireworks shows: A mix of lithium and strontium salts, along with some other chemicals, creates the show's brilliant red color.

Additional resources:

• Los Alamos National Laboratory: Lithium
• Jefferson Lab: The Element Lithium
• U.S. National Library of Medicine: Lithium (medication)

This article was updated on Oct. 23, 2018 by Live Science Managing Editor, Jeanna Bryner.

Diamonds are forever, and gold is precious, but which is rarer? And does that rarity have anything to do with the price we see at a jewelry store?

The answer, it turns out, isn't as "clear-cut" as you might think.

Gold — a heavy metal — is one of Earth's rarer elements, formed in the collisions of neutron stars, said Ulrich Faul, an Earth scientist and professor at the Massachusetts Institute of Technology.

Then, during the formation of Earth, the heaviest elements gravitated toward Earth's core, said Yana Fedortchouk, an Earth sciences professor and co-director of the Experimental High-Pressure Geological Research Laboratory at Dalhousie University in Halifax, in the Canadian province of Nova Scotia. That means that up near the Earth's crust, large amounts of gold are hard to find. [Photos Dazzling Minerals and Gems]

You can find it, though, in low concentrations. It's "present in a large variety of rocks in the crust," Fedortchouk told Live Science. "But in order to form a deposit, it needs to reach certain concentrations to make mining economically feasible."

According to Fedortchouk, the average concentration of gold in Earth's crust is "very, very low," at 4 parts per billion. In order to produce any minable concentration of gold that could be of market value, the gold deposit would have to be 1,250 times more concentrated, she said.

Diamonds, on the other hand, are a highly pressurized form of a very common element: carbon. In its non-pressurized form, it's known as graphite — the stuff in pencils. Compared with gold, the average concentration of carbon in Earth's crust is approximately 200,000 parts per billion, according to "Fluids in the Earth's Crust: Their Significance in Metamorphic, Tectonic, and Chemical Transport Processes" (Elsevier Science Ltd., 1978), a book written by the noted geologist William Fyfe, who died in 2013.

So, the rarity of diamonds has little to do with their elemental composition; rather, the natural transformation of carbon into diamonds that can be mined is an extremely arduous (and rarely successful) process.  

"Diamonds can only be produced in the Earth's mantle and somehow be brought to the surface, or they can be formed during meteorite impact," but those diamonds are small and never gems, Fedortchouk said. (The mantle is the layer of Earth beneath the crust.)"Diamonds formed deep in the Earth's mantle can be brought up by deep magma or pushed up during the slow uplift of deep rocks during mountain growth processes. But during slow uplift, diamonds get graphitized [turned into graphite] and never make it up to the surface as gem stones."

The formula required for diamonds to form depends on depth, temperature and pressure: Carbon is buried at least  93 miles (150 kilometers) beneath the Earth's surface, heated to about 2,200 degrees Fahrenheit (1,204 degrees Celsius) under approximately 725,000 pounds of pressure per square inch  (5 billion pascals), and then rapidly brought to the surface by a volcanic eruption to cool. This extraordinary process makes natural, minable diamonds rarer than gold, Fedortchouk said.

But, in its elemental form, gold is significantly rarer than diamonds, Faul told Live Science. After all, carbon is one of the most abundant elements on Earth — especially in comparison to heavier metals like gold — and diamond is simply composed of carbon under immense pressure. 

The invention of synthetic diamonds complicates the question even further. Scientists can re-create the conditions necessary to transform graphite into diamonds in a lab — no volcanic eruption necessary — but the same can't be said for gold (sadly, alchemy is still a pseudoscience). Even though synthetic diamonds are made of the same substance as natural diamonds, according to diamond designer Ritani, synthetic diamonds usually sell for 30 percent less on the market because they aren't considered as valuable.

But does the mere existence of lab-created diamonds make these gems more common than we thought? Faul argues that it does: "Diamonds below a certain size are not worth mining in the first place," he said. "Who wants to buy a diamond that needs a magnifying glass to be seen? Gold is more abundant than large diamonds, but diamonds as a class of material are not particularly rare. I think part of their reputation has to do with amazing public relations!"

Originally published on Live Science.

Radon (Rn) is a radioactive, colorless, odorless and tasteless gas that occurs naturally as the decay product of the elements radium, uranium and thorium. It is a noble (or inert) gas, meaning it is inactive chemically and combines with other substances only under extreme conditions. It is dense — the heaviest known gas — and it is considered a health hazard due to its radioactivity..

Radon is rare in nature because its isotopes are so short-lived and because its main source radium is also quite rare, according to Encyclopaedia Britannica. Radon has no known biological purpose but is believed to have played a major role in evolution, since radiation is required for genetic modifications to take place, according to the Royal Society of Chemistry. 

Just the facts

• Atomic number (number of protons in the nucleus): 86
• Atomic symbol (on the periodic table of the elements): Rn
• Atomic weight (average mass of the atom): 222
• Density: ;9.073 grams per liter
• Phase at room temperature: Gas
• Melting point: minus 95 degrees Fahrenheit (minus 71 degrees Celsius)
• Boiling point: minus 79 F (minus 61.7 C)
• Number of isotopes (atoms of the same element with a different number of neutrons): 3 naturally occurring (radon-219, radon-220, and radon-222); 33 whose half-lives are known with mass numbers 196 to 228; none are stable
• Most common isotope: Rn-222 (half-life of 3.823 days) 

Discovery

Credit for the discovery of radon is given to German chemist Friedrich Ernst Dorn in 1900, according to Chemicool. He discovered the new gas, which he referred to as radium "emanation," while studying radium's decay chain. Radium had been discovered just two years earlier by the Nobel Prize winning scientist Marie Curie.

Scottish chemist Sir William Ramsey, winner of the Nobel Prize for chemistry in 1904, explored the properties of radon even further. With the help of English chemist Robert Whytlaw Gray, Ramsey isolated radon and calculated its density so that it might be included correctly in the periodic table. They discovered that it was the heaviest gas ever known. They renamed the gas to "niton" after the Latin word for shining (nitens). But this name didn't stick either, and by 1923, it became known around the world as "radon," according to Chemicool.

Danger

Radon is present in the air nearly everywhere, and everyone breathes in radon every day, according to the National Cancer Institute. At low levels, it is harmless. However, people who inhale high levels of radon are at an increased risk of developing lung cancer. 

According to the Environmental Protection Agency (EPA), radon is the No. 1 cause of lung cancer among non-smokers and is associated with approximately 21,000 lung cancer deaths a year; 2,900 of those deaths occur among people who have never smoked.

Around one in 15 homes in the United States has elevated radon levels. The odorless gas can enter homes through cracks in walls, floors and foundations. It can also be released from building materials or through water originating from radon-contaminated wells, according to the National Institutes of Health (NIH). Radon levels can be greater in homes and buildings that are well-insulated, tightly sealed or built on soil rich in the radioactive elements uranium, radium and thorium.

A measurement of radioactivity is picocuries per liter of air (pCi/L). In the United States, the average indoor radon level is about 1.3 pCi/L, according to the EPA. The average outdoor level is about 0.4 pCi/L. The U.S. Surgeon General and EPA recommend fixing homes with radon levels at or above 4 pCi/L. The EPA also recommends that people think about fixing their homes for radon levels between 2 pCi/L and 4 pCi/L.

The 'Watras Incident'

In 1984, an odd coincidence known as the "Watras Incident" led to the discovery of the highest radon reading ever in Pennsylvania and ultimately urged the EPA to get involved in monitoring radon levels in residential homes.

Stanley J. Watras, a construction engineer at the Limerick nuclear power plant in Pottstown, Pennsylvania, set off the alarm at a radiation monitor installed to ensure that workers were not leaving the building with unsafe levels of radiation on their bodies.

This was quite a surprise, because the plant was still under construction and had not even been filled with nuclear fuel yet — so exposure would have been impossible. Eventually, a team of specialists discovered that Watras was not picking up the radiation at the plant, but from his own home —radiation levels in his home were 700 times higher than the maximum level considered safe.

The specialists discovered that the culprit was radon gas, which had been seeping into his home from underground, according to The Morning Call. Living there was the equivalent of smoking a couple hundred packs of cigarettes a day.

The family moved out immediately, and the home was turned into a scientific laboratory for the long-term measurement of radon and the testing of radon mitigation approaches. After several months, the radon was reduced to an acceptable level, and the family returned. Today, the U.S. Surgeon General and the EPA recommend that all homes be tested for radon.

Who knew?

• Radon was the fifth radioactive element to be discovered, after uranium, thorium, radium and polonium.
• Radon gas is colorless, but it exudes a brilliant yellow phosphorescence (light emitted from a substance without perceptible heat) at temperatures below its freezing point.
• Decades ago, radium salts were mixed into paints to make them glow in the dark. Once the EPA deemed radon a health risk, however, radon was pulled from consumer products.
• Hundreds of years ago, a wasting disease of miners was known as mala metallorum. In 1879, the condition was identified as lung cancer caused by exposure to radioactive substances, including uranium and radon.

Surprising research: Is radon beneficial in low doses?

For thousands of years, people have bathed in natural hot springs for the therapeutic properties of the water. Many of these spas have been found to contain radon. And yet,rather than developing radiation sickness or cancer, many claim that bathing in the radon-rich groundwater reduces inflammation and pain.

Also, radon was sometimes used by hospitals to treat cancer and other diseases. Hospitals used to produce it themselves by pumping radon from a radium source and sealing it in small tubes called seeds or needles. The seeds were injected at or near the site of the tumor. However, this is no longer a widespread practice.

A study led by researchers at Worcester Polytechnic Institute suggested that low levels of radon gas -- those levels typically found in 90 percent of American homes -- may actually reduce the risk of developing lung cancer by as much as 60 percent. The findings were published in the journal Health Physics in 2008. 

The results contrast the findings of previous studies, which suggested that low-level radon exposure is tied to a slightly elevated risk of lung cancer (with no statistical difference) or no risk at all, according to the study press release in Science Daily.

The study is the first to find a statistically significant hormetic effect of low-level radon exposure. The hormetic effect (hormesis) occurs when toxins and other environmental stressors have a beneficial effect at very low doses. The prevailing idea is that low doses of such toxins stimulate the immune system and the repair mechanisms in cells.

The researchers were quite surprised at the findings. In fact, the goal had been to establish what level of radon exposure was linked to lung cancer risk and to determine a safety zone for radon levels in the home, according to Science Daily. 

Ultimately, they found that the chance of developing lung cancer fell below one (the no effect level) at radon exposure within the range of 0-4 picoCuries per liter, according to Science Daily. This is the level of around 90 percent of homes in the United States. The EPA recommends that homeowners take action when exposure levels reach over 4 picoCuries per liter, due to the belief that increasing radon exposure is correlated to a progressively greater risk for cancer.

The researchers noted that the new findings do not dispute the lung cancer risk tied to greater levels of radon exposure, such as that experienced by uranium miners. However, the study shows a significant departure from previous results and beliefs about radon.

A 2011 research article published in the peer-reviewed journal Dose-Response showed similar results. The researchers reported that low-level residential radon is credited for an effect known as "activated natural protection" (ANP) against lung cancer, including smoking-related lung cancer.

Californium is a synthetic, radioactive element not found in nature. It is an actinide: one of 15 radioactive, metallic elements found at the bottom of the periodic table. The pure metal is silvery-white, malleable and so soft it can be easily sliced with a razor blade. Californium is moderately chemically reactive. It slowly tarnishes in air at room temperature — small pieces or foils of the metal begin to oxidize, but not violently.

Just the Facts

• Atomic number (number of protons in the nucleus): 98
• Atomic symbol (on the periodic table of the elements): Cf
• Atomic weight (average mass of the atom): 251
• Density: Unknown
• Phase at room temperature: Solid
• Melting point: 1,652 F (900 C)
• Boiling point: Unknown
• Number of isotopes (atoms of the same element with a different number of neutrons): 20 isotopes whose half-lives are known with mass numbers 237 to 256. 
• Most common isotopes: No naturally occurring isotopes

All californium isotopes are radioactive. Cf-251 is the most stable with a half-life of about 800 years. The isotope Cf-252 is a very strong neutron emitter, meaning it has the unusual property of giving off neutrons when it breaks apart. Isotopes that behave in this way are very unusual, according to Chemistry Explained. Cf-252 is used in a variety of industries: military applications, metal detectors, and as a way to detect metal fatigue and stress in airplanes. 

Discovery

In 1950, American scientists Stanley Thompson, Kenneth Street, Albert Ghiorso and Glenn Seaborg first produced californium in a lab at the University of California, Berkeley. It was the sixth synthetic transuranium ("beyond" uranium) element in the actinide group to be discovered. The discovery occurred when the chemists bombarded curium-242 with alpha particles (helium atoms without electrons) in a 60-inch cyclotron particle accelerator, according to Chemicool.com. 

Each nuclear reaction created Cf-245 — an isotope with a half-life of about 45 minutes — and a free neutron. The scientists produced around 700,000 atoms of Cf-245, just enough to make a cube with sides measuring only 27 nanometers, according to Chemicool. After a chemical analysis, the scientists confirmed that a new element had been discovered. 

In 1958, californium was isolated in larger quantities (1.2 micrograms) for the first time by Thompson and researcher Burris Cunningham at the Materials Testing Reactor in Idaho via prolonged (five years) neutron irradiation of plutonium-239. 

Uses

Cf-252, a californium isotope with a half-life of 2.645 years, is a very strong neutron source. One microgram (0.000001 grams) of Cf-252 produces 170 million neutrons per minute, according to Jefferson Lab. This isotope has a variety of uses, including acting as a neutron emitter, providing neutrons for the start-up of nuclear reactors. It is also used as a neutron source to detect gold and silver ores through a technique known as neutron activation. It is used in neutron moisture gauges, devices that can detect water and oil-bearing layers in oil wells, according to Jefferson Lab. 

The isotope Cf-252 is a target material for producing transcalifornium elements. For example, element 118, oganesson, the heaviest element, was produced when scientists bombarded californium with calcium ions, according to Chemicool. Cf-252 is also used to treat cervical cancer and can help analyze the sulfur content of petroleum.

New research

There is still so much to learn about the heaviest elements in the periodic table, such as californium, berkelium and einsteinium. Thomas E. Albrecht-Schmitt, a professor of chemistry at Florida State University and his research group study the chemical and physical properties of these elements. Their research into how they bond with other elements is rewriting textbooks.

One of the first problems of studying californium, however, is actually getting a piece of it. After years of negotiating with the U.S. Department of Energy, Albrecht-Schmitt was given 5 milligrams of californium through an endowment to the university. These 5 mg have stretched through several experiments, leading to the discovery that californium has the ability to bond with and separate other materials, according to a study press release in Science Daily. 

One of Albrecht-Schmitt's studies, published in the journal Nature Communications, revealed that californium is a transitional element, meaning that it helps link one part of the periodic table to the next. Specifically, californium shares properties with the three previous actinide elements in the table — americium, curium and berkelium — as well as the three actinides that come after it — einsteinium, fermium and mendelevium. 

The actinides consist of 15 radioactive, metallic elements, from actinium (Ac, atomic number 89) to lawrencium (Lr, atomic number 103) in the periodic table. The heavier members are extremely unstable and do not occur naturally.

"We discovered that californium represents a place in [the] series of actinide elements (Ac to Lr) where chemistry fundamentally changes with respect to the earlier actinides," Albrecht-Schmitt told Live Science.

"Starting at californium and strengthening through einsteinium, fermium and mendelevium, and finally maximizing at nobelium, the 2+ oxidation state starts to dominate the chemistry of these elements," he explained. "For californium through mendelevium, the 2+ state is metastable, and this creates chemical and physical properties that are significantly affected by charge transfer from the ligands to the metals."

"Physicists call this 'valence instability,' and it gives rise to a host of unanticipated electronic features such as bonding changes that are more similar to transition metals than they are to lanthanides."

Albrecht-Schmitt noted that this research is important for two reasons: "First, californium is the heaviest element for which bulk chemical and physical properties can be measured," he explains. "Its electronic properties are being significantly affected by its inner electrons moving at significant factions of the speed of light. Thus, californium is a test bed for measuring how relativistic effects alter chemistry."

"Second, while these properties are less pronounced in lighter actinides like plutonium, they are still present, and can be capitalized on in separation processes that are used to recycle used nuclear fuel and help mitigate the environmental legacy of the Cold War," he said.

Who knew?

• Californium is the heaviest element that has been produced in weighable amounts.
• The spectrum of Cf-254 has been observed in supernovae.
• The Oak Ridge National Laboratory in Tennessee is the only producer of Cf-252 for the U.S. government.
• Californium is considered to be a transuranium element, meaning "beyond uranium," on the periodic table. Uranium is element number 92, so elements with atomic numbers higher than 92 are transuranium elements.
• Cf-252 is used to treat some cervical and brain cancers when radiation is ineffective.
• Only a few compounds of californium have been produced and studied, including californium oxide (CfO3), californium trichloride (CfCl3) and californium oxychloride (CfOCl).

Is the sun due for a cosmic family reunion?

A new survey of 1 million stars in the Milky Way galaxy could help astronomers link our sun to its long-lost siblings.

The survey will identify stellar "DNA": the amounts of chemical elements — such as iron, aluminum and oxygen — that the stars contain. Astronomers could then use this data to find stars that emerged from the same birth clusters in galaxies' stellar nurseries, thereby matching stars to their "birth families," according to a statement released by The University of Sydney, one of several institutions participating in the astronomical survey. [What Will Happen to Earth When the Sun Dies?]

When the universe formed after the Big Bang, only two elements were present: hydrogen and helium. Elements that emerged later helped to shape stars and planets, making it possible for life to take hold on Earth. This new survey is measuring elements in more stars than in any previous project and at an unprecedented level of precision, which will help astronomers understand how galaxies form and change over time, university representatives said.

Today (April 18) marked the first data release from the enormous observational project, known as Galactic Archaeology with HERMES (GALAH). The project launched in 2013 and brings together astronomers from Europe and Australia, with the goal of observing more than 1 million stars. GALAH uses the HERMES instrument — its name stands for High Efficiency and Resolution Multi-Element Spectrograph — which is installed in the Anglo-Australian Telescope (AAT) in New South Wales, Australia. The instrument photographs light in four optical bands: red, blue, green and infrared, according to the Australian Astronomical Observatory (AAO).

In the data release, scientists described the observations of 340,000 Milky Way stars, reporting the findings in 11 studies released simultaneously in the journals Monthly Notices of the Royal Astronomical Society, and Astronomy and Astrophysics, according to the university's statement.

For the GALAH project, the AAT gathers starlight from 360 stars at once, and HERMES splits the light into spectra, or bands of light in different wavelength ranges. The size and placement of dark bands in the spectra reveal the amounts of different elements in a star, with each element emitting a unique signature pattern at different wavelengths, GALAH team member Daniel Zucker, an associate professor at Macquarie University in Australia, said in the statement.

Software nicknamed "the Cannon" (in honor of the pioneering U.S. astronomer Annie Jump Cannon) analyzes those bands in the spectra, looking for matches between stars. Earth's sun, like all stars, originated in a nursery cluster that likely spawned thousands of other stars. But because clusters in the Milky Way are usually swiftly ripped apart and scattered across the galaxy, it's difficult to say which stars in the galaxy were born in the same place, GALAH project scientist Sarah Martell, a senior lecturer at the University of New South Wales in Australia, reported in the statement. Gathering the stars' "DNA" and comparing "fingerprints" in spectra could help astronomers match the sun with the siblings that formed alongside it billions of years ago, according to Martell.

"No other survey has been able to measure as many elements for as many stars as GALAH," Gayandhi De Silva, a research astronomer with AAO and the HERMES instrument scientist overseeing GALAH's collaborators, said in the statement. "This data will enable such discoveries as the original star clusters of the galaxy, including the sun's birth cluster and solar siblings." 

Original article on Live Science.

Hafnium is a lustrous, silvery-gray transition metal. Discovered in 1923, it was the next-to-last element with stable nuclei to be added to the periodic table (the final one was rhenium in 1925). Hafnium is named after the Latin word for Copenhagen: Hafnia. The element has some very important commercial applications, including its use in the nuclear power industry, electronic equipment, ceramics, light bulbs and in the making of super-alloys.

Hafnium is rarely found free in nature, and instead is present in most zirconium minerals at a concentration of up to 5 percent. In fact, hafnium is so chemically similar to zirconium that separating the two elements is extremely difficult. Most commercial hafnium is produced as a byproduct of zirconium refining.

Hafnium is the 45th most abundant element on Earth, comprising about 3.3 parts per million (ppm) of the Earth's crust by weight, according to Chemicool. Hafnium is quite resistant to corrosion because of the formation of an oxide film on exposed surfaces. In fact, it is unaffected by water, air and all alkalis and acids except for hydrogen fluoride. 

Hafnium carbide (HfC) has the highest melting point of any known two-element compound at nearly 7,034 degrees Fahrenheit (3,890 degrees Celsius), according to Jefferson Lab. The compound hafnium nitride (HfN) also has a high melting point, around 5,981 degrees F (3,305 degrees C). Among compounds of three elements, the mixed carbide of tungsten and hafnium has the single highest melting point of any known compound at 7,457 degrees F (4,125 degrees C), according to Chemistry World. Some other hafnium compounds include hafnium fluoride (HfF₄) hafnium chloride (HfCl₄) and hafnium oxide (HfO₂).

Just the facts

• Atomic number (number of protons in the nucleus): 72
• Atomic symbol (on the periodic table of the elements): Hf
• Atomic weight (average mass of the atom): 178.49
• Density: 13.3 grams per cubic centimeter
• Phase at room temperature: Solid
• Melting point: 4,051 degrees Fahrenheit (2,233 degrees Celsius)
• Boiling point: 8,317 degrees F (4,603 degrees C)
• Number of isotopes (atoms of the same element with a different number of neutrons): 32 whose half-lives are known with mass numbers 154 to 185
• Most common isotopes: Hf-174, Hf-176, Hf-177, Hf-178, Hf-179 and Hf-180. 

Discovery

Hafnium's presence had been predicted decades before its discovery, according to Chemistry World. The element proved to be quite elusive, as it was nearly impossible to distinguish it chemically from the much more common zirconium. 

Hafnium was still unknown when Russian chemist and inventor Dimitri Mendeleev developed the Periodic Law — a pre-modern version of the periodic table of elements — in 1869. In his work, however, Mendeleev correctly predicted that there would be an element whose properties were similar to but heavier than zirconium and titanium.

In 1911, French chemist Georges Urbain, who had already discovered the rare earth element lutetium, believed he had finally discovered missing element 72 — which he proceeded to name celtium, according to Chemicool. However, a few years later his discovery was shown to be a combination of already discovered lanthanides (the 15 metallic elements with atomic numbers 57 through 71 in the periodic table).

It was still unclear whether missing element 72 would be a transition metal or a rare earth metal since it fell at the boundary between these two types of elements in the table. The chemists who believed it would be a rare earth element conducted many fruitless searches among minerals containing rare earths, according to Chemistry World.

However, new evidence arising from both the field of chemistry and physics supported the idea that element 72 would be a transition element. For example, scientists knew that element 72 fell below titanium and zirconium in the periodic table and both of these were known transition elements. In addition, Danish physicist Niels Bohr, one of the founders of quantum theory, predicted that element 72 would be a transition metal based on his electronic configuration for the element, according to Chemistry World.

In 1921, Bohr encouraged Hungarian chemist Georg von Hevesy and Dutch physicist Dirk Costerto — two young researchers in his institute at the time — to search for element 72 in zirconium ore. Based on his quantum theory of atomic structure, Bohr knew that the new metal would have a similar chemical structure to zirconium, so there was a strong chance the two elements would be found in the same ores, according to Chemicool.

Von Hevesy and Coster took Bohr's advice and proceeded to study zirconium ore using X-ray spectroscopy. They used Bohr's theory of how electrons fill shells and subshells within atoms to predict the differences between the two elements' X-ray spectra, according to Chemical and Engineering News. This method ultimately led to the discovery of hafnium in 1923. The discovery was one of the only six then remaining gaps in the periodic table. They named the new element after Bohr's hometown of Copenhagen (Hafniain Latin). 

Uses

Hafnium is remarkably corrosion-resistant and an excellent absorber of neutrons, allowing its use in nuclear submarines and nuclear reactor control rods, a critical technology used to maintain fission reactions. Control rods keep the fission chain reaction active but also prevent it from accelerating beyond control.

Hafnium is used in electronic equipment such as cathodes and capacitors, as well as in ceramics, photography flash bulbs and light bulb filaments. It is used in vacuum tubes as a getter, a substance that combines with and removes trace gases from the tubes, according to Jefferson Lab. Hafnium is commonly alloyed with other metals such as titanium, iron, niobium and tantalum. For example, heat-resistant hafnium-nobium alloys are used in aerospace applications, such as space rocket engines. 

The compound hafnium carbide has the highest melting point of any compound consisting of just two elements, allowing it to be used to line high-temperature furnaces and kilns, according to Chemicool. 

Who knew?

• Hafnium is pyrophoric (ignites spontaneously) in powder form.
• English chemist Henry Moseley was the scientist who realized that Georges Urbain's element "celtium" was not the true element located under zirconium. Unfortunately, World War I interrupted this young scientist's important research. Moseley dutifully enlisted in the Royal Engineers of the British Army and was killed by a sniper in 1915. His death led England to establish a new policy prohibiting prominent scientists from engaging in combat.
• In 1925, Dutch chemists Anton Eduard van Arkel and Jan Hendrik de Boer came up with a method for producing high-purity hafnium. To do this, the scientists decomposed hafnium tetraiodide on a hot tungsten wire resulting in a crystal bar of pure hafnium, according to Chemicool. This method is called the crystal bar process.
• The nuclear isomer of hafnium has long been debated as a potential weapon. In the Hafnium Controversy, scientists debate whether the element is capable of triggering a rapid release of energy.
• Although zirconium is chemically very similar to hafnium, it is unlike hafnium in that it is very poor at absorbing neutrons. Therefore zirconium is used in the outer layer of fuel rods where it is important that neutrons can travel easily.

Dating Earth's layers with hafnium

In a recent study, an international team of researchers was able to confirm that Earth's first crust formed around 4.5 billion years ago, thanks to their chemical analysis of hafnium in a rare meteorite. The researchers believe the meteorite originated from the asteroid Vesta, following a large impact that sent rock fragments to Earth, according to the study press release in Science Daily. According to the researchers, meteorites are pieces of the original materials that formed all planets. For the study, they measured the ratio of the isotopes hafnium-176 and hafnium-177 in the meteorite. This gave them a starting point for Earth's composition. They compared the results with the oldest rocks on Earth, essentially confirming that a crust had already formed on the surface of the Earth around 4.5 billion years ago. Their findings are published in the Proceedings of the National Academy of Sciences (PNAS).

It is said that we are all made of stardust, but stardust itself is formed from the very elements that are the building blocks of everything around us. From the fundamental elements such as oxygen and carbon — the most common elements in the human body — to elusive radioactive elements like francium, researchers are piecing together the secrets of the periodic table.

So whether it's a story on stars covered in unusual elements or a list of important elements you've never heard of, our expert writers and editors always have engaging news, features and articles about elements for you to read.

Discover more about elements

—Periodic Table of Elements

—How the Periodic Table of the Elements is arranged

—What is the universe made of?

Periodic table quiz

Paul Sutter is an astrophysicist at The Ohio State University and the chief scientist at COSI science center. Sutter is also host of Ask a Spaceman and Space Radio, and leads AstroTours around the world. Sutter contributed this article to Space.com's Expert Voices: Op-Ed & Insights.

Solid. Liquid. Gas. The materials that surround us in our normal, everyday world are divided into three neat camps. Heat up a solid cube of water (aka ice), and when it reaches a certain temperature, it changes phases into a liquid. Keep cranking the heat, and eventually, you'll have a gas: water vapor.

Every element and molecule has its own "phase diagram," a map of what you should expect to encounter if you apply a specific temperature and pressure to it. The diagram is unique to each element because it depends on the precise atomic/molecular arrangement and how it interacts with itself under various conditions, so it's up to scientists to tease out these diagrams through arduous experimentation and careful theory. [The Strangest Space Stories Of 2017]

When it comes to hydrogen, we usually don't encounter it at all, except when it's buddied up with oxygen to make the more familiar water. Even when we do get it by lonesome, its shyness prevents it from interacting with us alone — it pairs up as a diatomic molecule, almost always as a gas. If you trap some in a bottle and pull the temp down to 33 kelvins (minus 400 degrees Fahrenheit, or minus 240 degrees Celsius), hydrogen becomes a liquid, and at 14 K (minus 434 degrees F or minus 259 degrees C), it becomes a solid.

You would think that on the opposite end of the temperature scale, a hot gas of hydrogen would stay … a hot gas. And that's true, as long as the pressure is kept low. But the combination of high temperature and high pressure leads to some interesting behaviors.

Jovian deep dives

On Earth, as we've seen, hydrogen's behavior is straightforward. But Jupiter isn’'t Earth, and the hydrogen found in abundance within and beneath the great bands and swirling storms of its atmosphere can be pushed beyond its normal limits. 

Buried deep below the planet's visible surface, the pressures and temperature rise dramatically, and the gaseous hydrogen slowly gives way to a layer of supercritical gas-liquid hybrid. Due to these extreme conditions, the hydrogen can't settle into a recognizable state. It is too hot to stay a liquid but under too much pressure to float freely as a gas — it's a new state of matter.

Descend deeper, and it gets even stranger. 

Even in its hybrid state in a thin layer just beneath the cloud tops, hydrogen is still bouncing around as a two-for-one diatomic molecule. But at sufficient pressures (say, a million times more intense than the Earth's air pressure at sea level), even those fraternal bonds aren't strong enough to resist the overwhelming compressions, and they snap.

The result, below roughly 8,000 miles (13,000 km) under the cloud tops, is a chaotic mix of free hydrogen nuclei — which are just single protons — intermingled with liberated electrons. The substance reverts to a liquid phase, but what makes hydrogen hydrogen is now completely disassociated into its component parts. When this happens at very high temperatures and low pressures, we call this a plasma — the same stuff as the bulk of the sun or a lightning bolt.

But in the depths of Jupiter, the pressures force the hydrogen to behave much differently than a plasma. Instead, it takes on properties more akin to those of a metal. Hence: liquid metallic hydrogen.

That's metal

Most of the elements on the periodic table are metals: They're hard and shiny, and make for good electrical conductors. The elements get those properties from the arrangement they make with themselves at normal temperatures and pressures: They link up to form a lattice, and each donates one or more electrons to the community pot. These dissociated electrons roam freely, hopping from atom to atom as they please. 

If you take a bar of gold and melt it down, you still have all the electron-sharing benefits of a metal (except the hardness), so "liquid metal" isn't all that foreign a concept. And some elements that aren't normally metallic, like carbon, can take on those properties under certain arrangements or conditions. 

So, at first blush, "metallic hydrogen" shouldn't be that strange an idea: It's just a nonmetallic element that starts behaving as a metal at high temperatures and pressures. [Lab-Made 'Metallic Hydrogen' Could Revolutionize Rocket Fuel]

Once a degenerate, always a degenerate

What's the big fuss?

The big fuss is that metallic hydrogen is not a typical metal. Garden variety metals have that special lattice of ions embedded in a sea of free-floating electrons. But a stripped-down hydrogen atom is just a single proton, and there's nothing a proton can do to build a lattice. 

When you squeeze on a bar of metal, you're trying to force the interlocking ions closer together, which they absolutely hate. Electrostatic repulsion provides all the support a metal needs to be strong. But protons suspended in a fluid? That ought to be much easier to squish. How can the liquid metallic hydrogen inside Jupiter support the crushing weight of the atmosphere above it?

The answer is degeneracy pressure, a quantum mechanical quirk of matter under extreme conditions. Researchers thought conditions that extreme might be found only in exotic, ultradense environments like white dwarfs and neutron stars, but it turns out that we have an example right in our solar backyard. Even when electromagnetic forces are overwhelmed, identical particles like electrons can only be squeezed so tightly together — they refuse to share the same quantum mechanical state.

In other words, electrons will never share the same energy level, which means they will keep piling on top of each other, never getting closer, even if you squeeze really, really hard.

Another way to look at the situation is via the so-called Heisenberg uncertainty principle: If you try to pin down the position of an electron by pushing on it, its velocity can become very large, resulting in a pressure force that resists further squeezing.

So the interior of Jupiter is strange indeed — a soup of protons and electrons, heated to temperatures higher than that of the sun's surface, suffering pressures a million times stronger than those on Earth, and forced to reveal their true quantum natures. 

Learn more by listening to the episode "What in the world is metallic hydrogen?" on the Ask A Spaceman podcast, available on iTunes and on the web at askaspaceman.com. Thanks to Tom S., @Upguntha, Andres C., and Colin E. for the questions that led to this piece! Ask your own question on Twitter using #AskASpaceman or by following Paul@PaulMattSutterfacebook.com/PaulMattSutter.

Follow us @Spacedotcom, Facebook and Google+. Original article on Space.com. 

On Thanksgiving, enormous balloons of popular characters from cartoons, comics and animated TV shows and movies make their much-anticipated appearance in a stately procession down New York City avenues, part of the traditional Macy's Thanksgiving Day Parade, now in its 91st year.  

Helium is what keeps these balloons aloft — approximately 300,000 cubic feet (8,495 cubic meters) of helium, Live Science previously reported, which is roughly the same volume as three-and-a-half Olympic-size swimming pools.

Is this a frivolous use of a limited resource? Earth's helium supply is finite, and it has been dwindling for decades, with shortages reported in recent years. However, experts recently told Live Science that descriptions of helium shortfalls may have been — for lack of a better word — overinflated. [Beyond Balloons: 8 Unusual Facts about Helium]

Most of the helium on Earth originates in natural gas deposits — helium is produced as a by-product of natural gas — with the largest deposits found in the United States, Algeria and Qatar, John Hamak, a lead petroleum engineer with Helium Resources at the U.S. Bureau of Land Management (BLM), told Live Science.

The United States also maintains a Federal Helium Reserve, which was established in the 1960s, Hamak explained. For about a decade, the government acquired helium surpluses from private companies for storage in the reserve, eventually accruing so much that the program was halted. In 2005, the reserve began selling off its stored helium, Hamak said.

Lifting and cooling

Overall, the U.S. helium resources and reserves hold about 744 billion cubic feet (21 billion cubic meters) of the gas, according to a report published in January by the U.S. Geological Survey (USGS). Reserves and resources in the rest of the world account for an additional 1 trillion cubic feet (31 billion cubic meters) of helium, the USGS reported.

And how much of that helium is used up per year? In 2016, helium consumption in the U.S. was around 1.7 billion cubic feet (47 million cubic meters). According to the USGS report, only about 17 percent of that helium was designated as "lifting gas," which is made possible by helium's exceptionally low density. That percentage includes use for inflatables such as weather balloons, balloons for military scouting and balloons found in the Macy's parade.

Helium also has the lowest melting and boiling points of any element, and can create frigid environments, so it is frequently used in cryogenics, to chill fiber optics in computer network cables, and for cooling superconductors in magnetic resonance imaging (MRI) machinery, Hamak told Live Science. As helium does not react with anything, it can also create safe, inert environments for potentially volatile work, such as in underwater welding or for testing spaceflight fuel tanks, he said. [Why Does Helium Affect Your Voice?]

Supply and demand

Helium is a finite resource — it can't be reproduced chemically, and what we have on Earth is the result of billions of years of radioactive decay. This has raised concerns that wasteful, widespread use of the gas would quickly consume all the helium there is, and in 2010, reports grimly hinted that the world's supply of helium could be depleted as soon as 2035.

With more computers and smartphones produced each year, the demand for helium (which is used to protect hard drives and to manufacture semiconductor chips) is greater than ever — but is that enough to create a shortage? Not quite, Hamak said.

"Shortages from the early part of this decade were caused by a combination of events," Hamak explained. Political instability in the Middle East was partly responsible, interrupting the production and export of helium extracted from natural gas. At the same time, not enough helium was being processed in the U.S. to meet growing demands, prompting government officials to draw from the federal reserve more heavily as manufacturers scrambled to catch up, he said.

But within a few years, a new natural gas plant was up and running in Qatar, which not only took care of the helium shortage, it even initiated a brief period of oversupply, Hamak said.

Unfortunately, another shortage could happen anytime — there was a brief helium shortage this past summer, when Saudi Arabia and other countries blocked exports from Qatar — but the long-term prognosis for helium is still positive, he added.

"There are ample reserves in the U.S. and around the world, so I wouldn't say we'd ever run out, necessarily. It's just a matter of when a supply is disrupted — by a plant going down or by political disturbances," Hamak told Live Science.

And, it turns out, there's even more helium on Earth than was estimated a decade ago. Underground gas deposits found in 2016 in Tanzania's East African Rift Valley region were first thought to hold approximately 54 billion cubic feet (1.5 billion cubic meters) of helium. But that estimate was later revised to encompass a reserve close to twice that size, containing about 98.6 billion cubic feet (2.8 billion cubic meters) of helium gas.

Which means that the Thanksgiving balloon wranglers at Macy's won't need to retire their giant, inflatable versions of Charlie Brown, SpongeBob and Spider-Man just yet.

Original article on Live Science.

Bismuth is a brittle, crystalline, white metal with a slight pink tinge. It has a variety of uses, including cosmetics, alloys, fire extinguishers and ammunition. It is probably best known as the main ingredient in stomach ache remedies such as Pepto-Bismol.

Bismuth, element 83 on the periodic table of elements, is a post-transition metal, according to Los Alamos National Laboratory. (Different versions of the periodic table represent it as a transition metal.) Transition metals — the largest group of elements, which includes copper, lead, iron, zinc and gold — are very hard, with high melting points and boiling points. Post-transition metals share some characteristics of transition metals but are softer and conduct more poorly. In fact, bismuth's electric and thermal conductivity is unusually low for a metal. It also has a particularly low melting point, which enables it to form alloys that can be used for molds, fire detectors and fire extinguishers. 

Until recently, bismuth was considered the heaviest element that still had a stable nucleus. However, in 2003, researchers at the Institut d'Astrophysique Spatiale in Orsay, France, found that bismuth does decay into thallium, but it has an extremely long half-life: about 20 billion billion years (That's 20 followed by 18 zeroes). Put another way, if 100 grams of bismuth-209 (the natural isotope) had been present at the beginning of the universe more than 14 billion years ago, about 99.9999999 grams of it would still be around today, according to Chemicool. (Lead is now the heaviest stable element, according to Science magazine.)

Just the facts

• Atomic number (number of protons in the nucleus): 83
• Atomic symbol (on the periodic table of elements): Bi
• Atomic weight (average mass of the atom): 208.98040
• Density: 9.79 grams per cubic centimeter (5.6 ounces per cubic inch)
• Phase at room temperature: solid
• Melting point: 520.53 degrees Fahrenheit (271.4 degrees Celsius)
• Boiling point: 1,564 F (2,847 C)
• Number of natural isotopes (atoms of the same element with a different number of neutrons): 1
• Most common isotope: Bi-209

Discovery of bismuth

Though bismuth had been known as early as 1400, it was frequently confused with lead because it was similarly a heavy metal with a low melting point, according to the Royal Society of Chemistry. French chemist Claude Geoffroy the Younger was the first to prove that bismuth was distinct from lead in 1753.

The word "bismuth" is a Latinized version of an Old German word, "weissmuth" or "white substance," possibly named after the element's white oxide, according to Chemicool. 

Sources of bismuth

Naturally occurring bismuth is found in small quantities throughout Earth's crust both as a pure metal and combined with other elements in various compounds. The largest source of bismuth is found in the mineral bismuthinite, or bismuth sulfide (Bi₂S₃), according to John Emsley in his book, "Nature’s Building Blocks: An A-Z Guide to the Elements" (Oxford University Press, 1999). Bismuth is typically obtained as a by-product in refining lead, copper, tin, silver, and gold ores found in Bolivia, Peru, Japan, Mexico, and Canada.

Properties of bismuth

Compared to other metals, bismuth is the most diamagnetic; that is, it resists being magnetized and is repelled by a magnetic field, according to Chemicool. It also has low electric conductivity and the greatest electrical resistance when placed in a magnetic field, a trait called the Hall effect. 

It has a very low thermal conductivity — lower than any other metal except mercury, according to Chemicool. It also has a relatively low melting point, especially when alloyed with tin and lead. Bismuth burns with a blue flame and clouds of yellow oxide when heated in air, according to Robert E. Krebs in his book, "The History and Use of Our Earth’s Chemical Elements: A Reference Guide" (Greenwood Publishing Group, 2006). 

When liquid bismuth freezes, it expands rather than contracts because it forms a crystalline structure similar to water. Four other elements expand when they freeze: silicon, gallium, antimony and germanium, according to Chemicool. Bismuth crystals are easy to make at home. There are several how-to guides online, including this one at ScienceAlert and this one at ThoughCo.

Bismuth in your medicine cabinet

Bismuth subsalicylate (‎C₇H₅BiO₄), sold under the brand names Pepto-Bismol and Kaopectate, is a well-known remedy for stomach aches and diarrhea. It works by decreasing the flow of fluids and electrolytes into the bowel, reduces inflammation within the intestine, and may kill the organisms that can cause diarrhea, according to the U.S. National Library of Medicine. 

Many cosmetics including lipsticks, eyeshadow and nail polish contain bismuth oxychloride (BiOCl), a pearly powder that makes the products shine, according to Emsley.

Industrial and military uses of bismuth

According to Los Alamos:

• Bismuth compounds are used as catalysts in the manufacturing process of synthetic fiber and rubber. 
• When bismuth is combined with other metals such as lead, tin, iron and cadmium, it forms alloys with low melting points that can be used in fire detectors and extinguishers.
• Alloys of bismuth are also used in making sharp castings of objects subject to damage by high temperatures because the liquid metal expands 3.32 percent when it becomes solid.
• Bismanol, an alloy of bismuth and manganese, is a permanent magnet of high coercive force (a measure of magnetization) developed in the 1950s by the U.S. Naval Ordinance Laboratory in White Oak, Md. It was used in small motors. 
• The bismuth metal is used as replacement for lead in shot and bullets. 
• Bismuth can also be used in nuclear reactors and to make transuranium elements using a process called cold fusion. 

Additional resources

• Science magazine: Bismuth Not So Stable After All
• Here is more info on bismuth from the Los Alamos National Laboratory.
• Here is what the Jefferson Lab says about bismuth.
• The Royal Society of Chemistry also weighs in about bismuth. 

Neptunium, element 93 on the periodic table of elements, was the first transuranium element to be produced synthetically and the first actinide series transuranium element to be discovered. Its discovery came after several false findings of the element, including Enrico Fermi's attempt to bombard uranium with neutrons. That experiment resulted in the discovery of fission, or splitting atoms.

Neptunium is sandwiched on the periodic table between uranium and plutonium, which are also radioactive. All three of these elements, named after planets, have between 92 and 94 protons in their nuclei, large enough to undergo a nuclear fission reaction, or "atom splitting." Due to this capability, uranium and plutonium are both widely used in nuclear power plants and weapons. 

Neptunium, however, was discovered significantly later in history than either of its periodic table neighbors, and is not widely used. Neptunium remains an important element to study because it is produced by nuclear reactions of uranium and plutonium and can last as harmful radioactive waste for millions of years, according to a 2003 report by the Pacific Northwest Nuclear Laboratory. Understanding neptunium's chemistry is essential to ensure safe long-term nuclear waste storage. 

Just the facts

• Atomic number (number of protons in the nucleus): 93
• Atomic symbol (on the periodic table of elements): Np
• Atomic weight (average mass of the atom): 237
• Density: 11.48 ounces per cubic inch (19.86 grams per cubic cm)
• Phase at room temperature: solid
• Melting point: 1,191 degrees Fahrenheit (644 degrees Celsius)
• Boiling point: 7,052 F (3,900 C)
• Number of natural isotopes (atoms of the same element with a different number of neutrons): 4 — Neptunium-237 through neptunium-240. There are 21 additional known isotopes created in a lab.
• Most common isotope: Np-237

Discovery: Third time's a charm

According to John Emsley in his book, "Nature’s Building Blocks: An A-Z Guide to the Elements" (Oxford University Press, 1999), Italian scientist Enrico Fermi was the first to claim he discovered element 93, in 1934. He hypothesized that elements heavier than uranium (element 92) could be created by bombarding uranium with neutrons. Theoretically, this would add one neutral mass unit to the uranium atoms, which would then undergo beta decay, or loss of a negative charge that turns a neutron into a proton, resulting in an element with 93 total protons. Fermi's experiment did not end up producing an element; instead of the neutrons fusing with the uranium atoms, they split the uranium atoms into many fragment radioisotopes. Fermi was criticized for his false claim, and did not know at the time that he had actually performed the first atom splitting, or fission, experiment.

Just four years later in 1938, Romanian physicist Horia Hulubei and French chemist Yvette Cauchois made a similar false report of discovering element 93. They claimed that they found the element in a naturally occurring mineral sample. At the time, scientists rejected this, believing no elements with more protons than uranium (transuranium elements) were present in nature.

Element 93 was accepted as an existing element in 1940 at the University of California, Berkeley. Professor Edwin McMillan and graduate student Philip Abelson used a technique similar to Fermi, but with one important difference: they used slow-moving neutrons. McMillan used a machine called a cyclotron to slow the neutrons and then directed them at a uranium-238 target. This time, the neutrons actually worked to create element 93 by fusing with the uranium atoms instead of breaking them apart. Abelson analyzed the resulting sample, and noted unusual beta radiation that showed a new isotope (later named Np-289) was present. McMillan and Abelson decided to call the element neptunium because Neptune is the next planet beyond Uranus in the solar system. The discovery was the first transuranium element to be synthesized in a lab and earned McMillan a Nobel Prize in 1951.

Sources of neptunium

Although scientists used to think neptunium could only be created synthetically, trace amounts of four of neptunium's 25 isotopes have since been found in nature, according to Los Alamos National Laboratory. Uranium, found in rock, soil and water, undergoes a natural nuclear reaction that results in small amounts of isotopes Np-237 to Np-240. 

The majority of neptunium, however, is anthropogenic; that is, it is created as a byproduct of reactions in nuclear power plants. Scientists can extract neptunium from spent nuclear fuel in large quantities. Because of its long half-life of 2.14 million years, Np-237 is the most abundant isotope of neptunium created. Most other isotopes of neptunium have short half-lives and decay within days.

Properties of neptunium

Neptunium is a member of the actinide series, row 5f of the periodic table. This row (along with the lanthanide row above) is often depicted below and separately from the rest of the periodic table because it is too long to fit on a page with normal dimensions. All 15 actinide elements have very large atomic radii and are radioactive.

Neptunium is a silver metal and is very reactive, with four different oxidation states. When it combines with other elements it occurs as different colored solutions (purple, yellow, green and pink). Even on its own, neptunium occurs as three different allotropes, or physical forms, depending on the temperature. It is the densest of the actinides and can remain a liquid for the largest temperature range of any known element.

Can we use it?

Neptunium's current applications are limited. Neptunium has only been considered, not actually used, as a fissile nuclear fuel. However, neptunium-237 is used to create plutonium-238, which is then used in special energy generators that can power satellites, spacecraft and lighthouses for a long period of time. Neptunium-237 is also used in nuclear physics research as a part of a device that detects high-energy neutrons.

Can it harm us?

There may be radioactive neptunium in your house! Neptunium accumulates in a common household item: ionizing smoke detectors. To detect smoke, another actinide element, americium-241, emits radiation and turns into neptunium-237. But no need to worry: the amount of radioactive material in smoke detectors is negligible and causes no harm to human health, according to Emsley. Smoke detectors contain less than 0.00000001 ounces (0.0000003 grams) of americium, which decays so slowly that only about 0.2 percent of this already small amount converts to neptunium each year. 

Scientists are, however, concerned with the long-term storage of neptunium present in spent nuclear fuel, according to a 2005 article published by Berkeley Lab. Although neptunium makes up only a small percentage of total radioactive waste, it poses a particular threat because it is long-lasting and hard to extract. Amy Hixon, an assistant professor at the Notre Dame College of Engineering, has studied the less familiar actinide elements and how to best contain them. 

"The neptunium present in a used nuclear fuel rod can last for millions of years, and I'm not exaggerating," Hixon said as she explained the reality of containing neptunium. Her lab studies how neptunium and other actinides move through materials simulating geological repositories, like the one proposed for Yucca Mountain in Nevada. Though these deep storage sites are generally accepted as the safest long-term storage, there are none currently operating in the United States. The Yucca Mountain Nuclear Waste Repository was defunded under the Obama administration in 2011. The Trump administration has cut all funding for deep borehole waste research, but Congress may reconsider funding in the next budget cycle for 2018.

Additional resources

• Here is more info on neptunium from the Los Alamos National Laboratory.
• Here is what the Jefferson Lab says about neptunium.
• The Royal Society of Chemistry also weighs in about neptunium. 

The catchphrase "bazinga" — a zinger commonly uttered by Dr. Sheldon Cooper, a fictional theoretical physicist on the TV show "The Big Bang Theory" — has inspired the creation of a novel compound, according to new research.

The compound, known as BaZnGa (barium, zinc and gallium), may have unorthodox roots (after all, how many compounds exist because of a sitcom wisecrack?), but its creation fits into a goal pursued by countless physicists: to discover new materials, as well as materials with novel properties, said study co-author Paul Canfield, a professor of physics and astronomy at Iowa State University.

"It's part of our effort to try to discover new materials and characterize them," Canfield told Live Science. [Elementary, My Dear: 8 Little-Known Elements]

The idea to create BaZnGa came from study co-author Na Hyun Jo, a graduate student of physics in Canfield's lab. One evening, she turned on the TV and saw a commercial for "The Big Bang Theory" that showed Sheldon's "Bazinga" as the elements barium, zinc and gallium.

It just so happened that Jo and Canfield were already planning to study three-part compounds (known as a ternary) that contained barium and zinc. When Jo saw "BaZnGa" flash on the TV, it seemed to be a message from the universe, telling her, "'Come on guys, add some gallium, see what you can find!'" Canfield said.

The physicists first checked whether anyone else had ever made BaZnGa. The coast was clear, they found: There was no other reported compound containing just these three elements, Canfield said.

So, Canfield, Jo and their colleagues went ahead and made BaZnGa. The results were surprising: the compound forms a new, never-before-seen crystal structure. This finding is useful, because "whenever you find a new structure, it's nice because it gives you further information on how nature arranges atoms," Canfield said.

However, BaZnGa didn't have any other exciting features — at least any that they could find. Rather, it's like other nonmagnetic quasicrystals, the researchers said. (A quasicrystal shares some qualities with crystals, such as having diffraction, but its atoms aren't arranged as regularly as they are in real crystals, Live Science reported previously.)

But Canfield and Jo did find an amusing flip side to their new concoction: Sheldon says "bazinga" when he's attempting to make a jest or jape, so it's almost as if Sheldon is playing a joke on them, Canfield said.

"We were hoping that we would have a title, 'New high-temperature superconductor BaZnGa,'" Canfield said. "Not so."

For instance, "We checked to see if this new compound was quasicrystalline," the researchers wrote in the study — but it wasn't. The same thing happened when the researchers tested it for high-temperature superconductivity, as well as a barrage of other measures.

Nope. Bazinga again.

"In conclusion, Dr. Sheldon Cooper continues to be both a motivation and an enigma," the researchers wrote in the study. "We admire the accuracy of the prediction of BaZnGa, but regret having fallen prey to his legendary wit; BaZnGa!"

The scientists wrote a lighthearted and a serious version of the study. Both are available on ArXiv, a site where research that has yet to be peer-reviewed can be posted.

Canfield plans to present the results at the annual meeting of the American Physical Society in Los Angeles in March 2018, and he hopes that actor Jim Parsons, who plays Sheldon Cooper, or any of "The Big Bang Theory" cast members can attend "and perhaps discuss BaZnGa in the light of our findings," Canfield said.

Original article on Live Science.

Astatine is the rarest element on Earth; only approximately 25 grams occur naturally on the planet at any given time. Its existence was predicted in the 1800s, but was finally discovered about 70 years later. Decades after its discovery, very little is known about astatine. Indeed, physicists infer many of its properties — such as its radioactive properties, conduction and color — based on other halogen group members.

History

Dmitri Mendeleyev, the Russian chemist who in 1869 organized the elements into the periodic table that is still used today, predicted properties of the unknown element that would fill the blank space on the periodic table for element No. 85, according to Peter van der Krogt, a Dutch historian. Mendeleyev named this unknown element eka-iodine due to its position directly below iodine in the halogen group of elements. 

As the search for the new element began, several reports were published about element 85, according to a 2010 article published in the Bulletin for the History of Chemistry by Brett F. Thornton and Shawn C. Burdette, researchers in Sweden and the United States, respectively. These reports included claims that the element couldn't exist, that researchers finding the element were unable to isolate it, and that the reported properties were inconsistent with tests.

There is a great deal of ambiguity as to who first discovered astatine, according to Thornton and Burdette. The discovery could be attributed to a handful of researchers, most notably one of the following groups. 

The first claim that the mystery element had been discovered was in 1931 by Fred Allison at the Alabama Polytechnic Institute, according to Thornton and Burdette. Allison suggested the name "alabamine" for the new radioactive element that he had discovered. However, as no other researchers were able to replicate his results, and because several faults were found in his equipment, the search for the elusive element continued. Not before, however, the discovery was published in a few student textbooks.

Horia Hulubei and Yvetter Cauchois, researchers at the Sorbonne in Paris, published the results of their discovery of element 85 in 1938. They used chemical separation and published that they found three X-ray spectral lines for the element that closely matched previous predictions. Unfortunately, the breakout of World War II disrupted their research as well as communications among scientists around the world.

The first successfully recognized discovery of astatine was in 1940 by Dale R. Coson, Kenneth Ross Mackenzie and Emilio Segrè, researchers at the University of California Berkeley, according to Chemicool. As no one had been able to find the rare element in nature, this group of scientists artificially produced it by bombarding bismuth-209 with alpha particles in a particle accelerator. This reaction created astatine-211 as well as two free neutrons. The element was highly radioactive and unstable, which led to the name astatine from the Greek word that meant "unstable."

Yet another group of researchers independently identified and characterized element 85 in the early 1940s, according to Thornton and Burdette. Berta Karlik and Traude Bernert in 1942 reported the results of their studies, including the proposed name "viennium." However, because of WWII, the news was kept inside the German territories, and science news from other regions of the world was not brought in, so Karlik and Bernert were not aware of similar results from the Berkeley group. When Karlik and Bernert were made aware of the published results from the group at Berkeley, they still continued to study element 85 and added greatly to the knowledge about the decay chain that forms the element. 

Just the facts

• Atomic number (number of protons in the nucleus): 85
• Atomic symbol (on the periodic table of elements): At
• Atomic weight (average mass of the atom): 210
• Density: approximately 4 ounces per cubic inch (approximately 7 grams per cubic cm)
• Phase at room temperature: solid
• Melting point: 576 degrees Fahrenheit (302 degrees Celsius)
• Boiling point: unknown
• Number of natural isotopes (atoms of the same element with a different number of neutrons): at least 30 radioactive isotopes
• Most common isotopes: At-210 (negligible percent of natural abundance), Am-211 (negligible percent of natural abundance)

Who knew?

• Astatine is named after the Greek word 'astatos,' which means unstable, according to the Jefferson Laboratory.
• There are only about 25 grams of naturally occurring astatine in Earth's crust at any given time, according to Chemicool.
• According to Lenntech, astatine is the heaviest known halogen. According to Elemental Matter, halogen elements, including astatine, share similar properties; they are non-metals, have low melting and boiling points, are brittle when solid, are poor conductors of heat and electricity, and are diatomic (their molecules contain two atoms).
• Astatine is the least reactive and has the most metallic properties of any element in the halogen group, according to Chemicool.
• The isotope of astatine with the longest half-life is astatine-210 with a half-life of 8.1 hours, according to the Jefferson Laboratory.
• Many physical properties of astatine are still unknown, including its color, according to a 2013 article by D. Scott Wilbur published in Nature. Based on the color patterns shown by other members of the halogen family, it is believed that astatine is dark, probably close to black.
• Astatine is highly radioactive yet poses nearly no health or environmental effects at all due to its rarity and very short half-lives, according to Lenntech. Although if one does come into contact with it, astatine is thought to accumulate in the thyroid gland similarly to iodine.

Current research

The scarcity of astatine makes it an incredibly difficult element to study. Nevertheless, some researchers think astatine may have uses in treating cancer. Astatine may behave like iodine, which tends to collect in the thyroid gland, according to Chemistry Explained. Astatine may also go to the thyroid, and its radiation could kill cancer cells in the gland.

In a 2015 paper published in the International Journal of Molecular Sciences, a group of French researchers led by Françoise Kraeber-Bodéré describe a radioimmunotherapy (RIT) method of cancer therapy that uses radionuclides that emit either beta or alpha particles. Astatine-211 is one such isotope that could be beneficial to alpha therapy because it has a longer half-life than traditionally used bismuth-213, and it can be produced in particle accelerators. Astatine-211 has been studied for this use since at least 1989, according to the authors, and has shown to have promising results, including trials with bone marrow transplants in leukemia patients, stem cell transplantation studies in mice, and in chemotherapy treatments with patients with brain tumors.

The conclusions reached by the researchers shows that using a radioactive isotope, such as astatine-211, can improve RIT efficiency for treating tumors and other cancers, especially if treatment is started early in the disease. This method of RIT also has the potential to kill remaining tumor cells that are typically resistant to chemo and radioactive therapy.

Additional resources

• Chemicool: Astatine Element Facts
• Royal Society of Chemistry: Astatine
• Nature: Enigmatic Chemistry

Americium, a silvery-white, synthetic element, is created during nuclear reactions of heavy elements. The element and its isotopes have very few but important uses including smoke detectors found in nearly all buildings and the potential to power future space missions.

Americium is a highly radioactive element that can be dangerous when handled incorrectly and can cause severe illnesses. Since it is not found naturally in the environment, there is very little chance that humans and animals would be affected by the element unless they are in very close proximity to plutonium-based nuclear reactors.

Just the facts

• Atomic number (number of protons in the nucleus): 95
• Atomic symbol (on the periodic table of elements): Am
• Atomic weight (average mass of the atom): 243
• Density: 7.91 ounces per cubic inch (13.69 grams per cubic cm)
• Phase at room temperature: solid
• Melting point: 2,149 degrees Fahrenheit (1,176 degrees Celsius)
• Boiling point: 3,652 F (2,011 C)
• Number of natural isotopes (atoms of the same element with a different number of neutrons): 0. There are at least 19 radioactive isotopes created in a lab.
• Most common isotopes: Am-241 (negligible percent of natural abundance), Am-243 (negligible percent of natural abundance)

History

Glenn Seaborg, Albert Ghiorso, Ralph James and Tom Morgan discovered americium, as well as curium, in 1944 during their work at the wartime Metallurgical Laboratory at the University of Chicago (now known as Argonne National Laboratory), according to a 2008 article in the Bulletin for the History of Chemistry by Keith Costecka, an American chemist and environmental scientist. The researchers produced the synthetic element by bombarding plutonium-239 with neutrons to create plutonium-240, and then again to create plutonium-241. The plutonium-241 then decayed to americium-241. Americium is the third synthetic transuranic element and the fourth to be discovered.

The discoveries of americium and curium were announced in late 1945 by Seaborg on the live radio show Quiz Kids, according to a 2017 Nature article by Ben Still, a British scientist and author. The announcement was meant to have occurred five days later at a national meeting of the American Chemical Society. The element was named by researchers for the country that discovered it as well as a mirror to neighboring lanthanide element number, europium.

Americium was very difficult to isolate from curium and the process took over a year, according to Peter van der Krogt, a Dutch historian. The researchers nicknamed the elements pandemonium and delirium and even suggested that those names should become the elements' official names. Despite the researchers suggestions, the elements were given the names americium, after the continent of discovery and following the example of europium, and curium named for scientists Marie and Pierre Curie. The first substantial enough amount of americium that could be visibly studied was created in 1951, according to the Los Alamos National Laboratory. 

Who knew?

• According to 1986 article published in Radiochemistry and Nuclear Chemistry by Norman Edelstein and Lester Morss, American researchers, americium is one of 15 actinide metals. Actinide elements range from atomic numbers 89 (actinium) through 103 (lawrencium). These elements are all radioactive with an unusual range of physical properties.
• The primarily isotopes are americium-241 and americium-243, which have half-lives of about 433 years and 7370 years, respectively, according to the Royal Society of Chemistry.
• According to Lenntech, americium most likely occurs naturally in incredibly trace amounts in uranium minerals due to nuclear reactions. Past concentrations of americium may have been higher when local concentrations of uranium were higher and produced more nuclear reactions.
• Americium is primarily produced in nuclear reactors via the bombardment of plutonium with neutrons, according to the Royal Society of Chemistry.
• Americium is a portable source of gamma-rays and alpha particles for a variety of medical and industrial uses such as radiography and spectroscopy, helping to create flat glass by gauging its thickness. 
• According to the World Nuclear Association, smoke detectors that use americium are popular in homes and are sensitive to the presence of smoke or heat. These smoke detectors are relatively inexpensive and are sensitive to a wide range of fire conditions. The isotope americium-241 is used in these detectors as americium dioxide (AmO2).
• The Centers for Disease Control and Prevention states that americium-241 dust can cause certain cancers and can be swallowed, absorbed through a wound, or inhaled. The element tends to concentrate itself in the bones, liver, and muscles. Due to the longevity of the isotope, americium-241 can stay in the body for decades. 

Current research

Due to its rarity and radioactivity, the uses for americium are few. One such use for americium that is currently being researched is in batteries, specifically "space batteries." Research conducted by the United Kingdom's National Nuclear Laboratory (NNL), in conjunction with the European Space Agency (ESA), has shown promising results for obtaining the materials needed to build the americium-241 powered batteries. Researchers at the NNL were successfully able to design a method and successfully isolate americium-241 from plutonium. Further studies are currently being completed to examine the impacts that a larger-scale americium processing plant would have on the environment, as well as how to keep workers at such a plant safe. The long-term plan put forth by NNL would work to create a larger quantity of americium that can be used in the batteries. 

In a 2008 article presented at a European Space Power Conference, K. Stephenson and T. Blancquaert, scientists at ESA based in the Netherlands, said that plutonium has been the favored fuel source due to its high power output and 88-year half-life. The isotope of plutonium that is required for such space missions is very expensive with an extremely limited supply and with restrictive regulations. Americium-241, on the other hand, only has about a quarter of the power output that plutonium has, but it has a longer half-life (433 years), more easily produced, and can potentially bring cost and weight down by about a third.

Another group of scientists in Israel is conducting tests on batteries powered by americium-242, according to a 2008 article presented at the Congress of Nuclear Societies by M. Kurtzhand, et al., a group of Israeli nuclear engineers. The researchers said that americium-242 has a high power output, and, according to an article published on The Future of Things about the project, could power the International Space Station for up to 80 days. The battery powered by americium-242 faces some difficulties due to americium-242 being more difficult to produce than americium-241, but the isotope leads to ideal battery properties such as the ability to be made incredibly thin (about one micron) and with no moving parts leading to a robust and reliable power source.

This story was updated April 4 at 12:24 p.m. EDT.

Hydrogen is the most common element in the universe, but why is that?

To answer this question, "we need to go back to the Big Bang," said May Nyman, a professor of chemistry at Oregon State University.

The Big Bang created the elements on the periodic table, building blocks that help make up the universe. Each element has a unique number of subatomic particles: protons (positively charged), neutrons (neutral) and electrons (negatively charged). [What Are the Ingredients of Life?]

Hydrogen — with just one proton and one electron (it's the only element without a neutron) — is the simplest element in the universe, which explains why it's also the most abundant, Nyman said. (However, an isotope of hydrogen, called deuterium, contains one protron and one neutron, and another, known as tritium, has one proton and two neutrons.)

In stars, hydrogen atoms fuse to create helium — the second most common element in the universe, according to Encyclopedia.com. Helium has two protons, two neutrons and two electrons. Together, helium and hydrogen make up 99.9 percent of known matter in the universe, according to Encyclopedia.com.

Even so, there is still about 10 times more hydrogen than helium in the universe, Nyman said. Oxygen, the third most common element, is about 1,000 times less abundant than hydrogen, she added.

In general, the higher the atomic number, the less abundant is the element is, Nyman said.

Earth's composition, however, is different from that of the entire universe. For instance, oxygen is the most common element by weight in Earth's crust, followed by silicon, aluminum and iron, according to HyperPhysics, a site run by Georgia State University.

In the human body, the most common element by weight is oxygen, followed by carbon and hydrogen, according to HyperPhysics.

Hydrogen has a number of key roles in the human body. Hydrogen bonds help give DNA its signature twist, and it helps the stomach and other organs maintain the correct pH, or how acidic or basic it is, Nyman said.

"If your stomach gets too basic, hydrogen will be released to what it's bonded to," she said. "If it's too acidic, [hydrogen] will bond to something."

In addition, hydrogen allows ice to float on water (H20) because the hydrogen bonds push the frozen water molecules apart, making them less dense.

"Usually, substances are more dense when they're solid than when they're liquid," Nyman said. "Water is the only substance that is less dense than when it's [a] solid."

However, hydrogen can also be dangerous. Hydrogen gas reacting with oxygen led to the Hindenburg blimp catastrophe that killed 36 people in 1937, according to Airships.net. Moreover, hydrogen bombs can be incredibly destructive, although they have never been used as a weapon, "just demonstrated by the United States, USSR, Great Britain, France and China in the 1950s," Nyman said.

Hydrogen bombs, like atomic bombs, use a combination of nuclear fusion and fission reactions to cause destruction, and release both radiation and mechanical shock waves when detonated, she said. 

Editor's Note: This article has been updated to include information about hydrogen isotopes and to say that H-bombs can release radiation.

Original article on Live Science.

Named for the Norse god of thunder, thorium is a silvery, lustrous and radioactive element with potential as an alternative to uranium in fueling nuclear reactors. 

Just the facts

• Atomic number (number of protons in the nucleus): 90
• Atomic symbol (on the Periodic Table of Elements): Th
• Atomic weight (average mass of the atom): 232.0
• Density: 6.8 ounces per cubic inch (11.7 grams per cubic cm)
• Phase at room temperature: Solid
• Melting point: 3,182 degrees Fahrenheit (1,750 degrees Celsius)
• Boiling point: 8,654 F (4,790 C)
• Number of natural isotopes (atoms of the same element with a different number of neutrons): 1. There are also at least 8 radioactive isotopes created in a lab.
• Most common isotopes: Th-232 (100 percent of natural abundance)

History

In 1815, Jöns Jakob Berzelius, a Swedish chemist, first thought he had discovered a new Earth element, which he named thorium after Thor, the Norse god of war, according to Peter van der Krogt, a Dutch historian. In 1824, however, it was determined that the mineral was in fact yttrium phosphate.;

In 1828, Berzelius received a sample of a black mineral found on Løvø Island off the coast of Norway by Hans Morten Thrane Esmark, a Norwegian mineralogist. The mineral contained nearly 60 percent of an unknown element, which took over the name thorium; the mineral was named thorite. The mineral also contained many known elements, including iron, manganese, lead, tin, and uranium, according to Chemicool.

Berzelius isolated thorium by first mixing thorium oxide found in the mineral with carbon to create thorium chloride, which was then reacted with potassium to yield thorium and potassium chloride, according to Chemicool.

Gerhard Schmidt, a German chemist, and Marie Curie, a Polish physicist, independently discovered that thorium was radioactive in 1898 within a couple months of each other, according to Chemicool. Schmidt is often credited with the discovery.

Ernest Rutherford, a New Zealand physicist, and Frederick Soddy, an English chemist, discovered that thorium decays at a fixed rate into other elements, also known as the half-life of an element, according to Los Alamos National Laboratory. This work was critical in furthering the understanding of other radioactive elements.

Anton Eduard van Arkel and Jan Handrik de Boer, both Dutch chemists, isolated high purity metalic thorium in 1925, according to Los Alamos National Laboratory. 

Who knew?

• In its liquid state, thorium has a greater temperature range than any other element, with nearly 5,500 degrees Fahrenheit (3,000 degrees Celsius) between melting and boiling points, according to Chemicool.
• Thorium dioxide has the highest melting point of all known oxides, according to Chemicool.
• Thorium is about as abundant as lead and at least three times as abundant as uranium, according to Lenntech.
• The abundance of thorium in Earth's crust is 6 parts per million by weight, according to Chemicool. According to Periodic Table, thorium is the 41st most abundant element in Earth's crust.
• Thorium is mainly mined in Australia, Canada, the United States, Russia and India, according to Minerals Education Coalition.
• Trace levels of thorium are found in rocks, soil, water, plants and animals, according to the U.S. Environmental Protection Agency (EPA).
• Higher concentrations of thorium are typically found in minerals such as thorite, thorianite, monazite, allanite, and zircon, according to Los Alamos National Laboratory.
• The most stable isotope of thorium, Th-232, has a half-life of 14 billion years, according to the EPA.
• According to Los Alamos, thorium is created in the cores of supernovae and then scattered across the galaxy during the explosions.
• Thorium had been used since 1885 in gas mantles, which provide the light in gas lamps, according to Los Alamos. Due to its radioactivity, the element has been replaced by other nonradioactive rare-earth elements.
• Thorium is also used to strengthen magnesium, coating tungsten wire in electrical equipment, controlling the grain size of tungsten in electric lamps, high-temperature crucibles, in glasses, in camera and scientific instrument lenses, and is a source of nuclear power, according to Los Alamos.
• Other uses for thorium include heat-resistant ceramics, aircraft engines, and in light bulbs, according to Chemicool.
• According to Lenntech, thorium was used in toothpaste until radioactivity dangers were discovered.
• Thorium and uranium are involved in the heating of Earth's interior, according to Minerals Education Coalition.
• Too much thorium exposure can cause lead to lung disease, lung and pancreatic cancer, alter genetics, liver disease, bone cancer, and metal poisoning, according to Lenntech.

Current research

A great deal of research is going into using thorium as a nuclear fuel. According to an article from the Royal Society of Chemistry, thorium used in nuclear reactors provide many benefits over using uranium:

• Thorium is three to four times more abundant than uranium.
• Thorium is more easily extracted than uranium.
• Liquid fluoride thorium reactors (LFTR) have very little waste compared with reactors powered by uranium.
• LFTRs run at atmospheric pressure instead of 150 to 160 times atmospheric pressure currently needed.
• Thorium is less radioactive than uranium.

According to a 2009 paper by NASA researchers Albert J. Juhasz, Richard A. Rarick, and Rajmohan Rangarajan, thorium reactors were developed at the Oak Ridge National Laboratory in the 1950s under the direction of Alvin Weinberg for supporting nuclear aircraft programs. The program stopped in 1961 in favor of other technologies. According to the Royal Society of Chemistry, thorium reactors were abandoned because they did not produce as much plutonium as uranium-powered reactors. At that time, weapons-grade plutonium, as well as uranium, was a hot commodity due to the Cold War.

Thorium itself is not used for nuclear fuel, but it is used to create the artificial uranium isotope uranium-233, according to the NASA report. Thorium-232 first absorbs a neutron, creating thorium-233, which decays to protactium-233 over the course of about four hours. Protactium-233 slowly decays to uranium-233 over the course of about ten months. Uranium-233 is then used in nuclear reactors as fuel.

Additional resources

• Bulletin of the Atomic Scientists: Thorium: The Winder Fuel That Wasn't
• Jefferson Lab: The Element Thorium
• USGS: Thorium Statistics and Information

Einsteinium, the 99th element on the Periodic Table of Elements, is a synthetic element that is produced in extremely small amounts and with a very short lifetime. If the name seems familiar, it's because it is indeed named after famed physicist Albert Einstein, although he had nothing to do with the element's discovery or research.

History

Einsteinium was discovered during the examination of debris from the first hydrogen bomb test in November 1952, according to Chemicool. A team of scientists from the Lawrence Berkeley National Laboratory, the Argonne National Laboratory and the Los Alamos Scientific Laboratory and led by Albert Ghiorso, an American nuclear scientist at Berkeley, studied the debris collected by drones using chemical analysis. Minuscule amounts of einsteinium-253, an isotope of einsteinium, were discovered (less than 200 atoms, according to an article printed in Nature Chemistry by Joanne Redfern, a British science writer, in 2016). Fermium, the 100th element, was also discovered in the debris.

The results of the test were not published until 1955, according to Peter van der Krogt, a Dutch historian. At the time of the hydrogen bomb demonstration, tensions due to the Cold War were running high and many new discoveries were kept secret. Due to the method of creation and the nature of the new elements, further research of einsteinium continued in silence, according to the Los Alamos National Laboratory.

Further production and research on einsteinium, as well as fermium, was done at the Oak Ridge National Laboratory in Tennessee, according to Lenntech. The scientists created larger amounts of the element and also developed methods of purifying the einsteinium, as described in a 1978 paper by D.E. Ferguson and presented at a symposium. The paper describes the production of einsteinium and fermium in nuclear reactors by bombarding heavy elements such as uranium and curium with neutrons and having those products undergo radioactive decay. Einsteinium and the other heavy elements were then extracted from a tank filled with a solvent.

Just the facts

• Atomic number (number of protons in the nucleus): 99
• Atomic symbol (on the Periodic Table of Elements): Es
• Atomic weight (average mass of the atom): 252
• Density: Unknown
• Phase at room temperature: solid
• Melting point: 1,580 degrees Fahrenheit (860 degrees Celsius)
• Boiling point: Unknown
• Number of natural isotopes (atoms of the same element with a different number of neutrons): 0. There are also at least 19 artificial isotopes created in a lab, none of which are stable.

Who knew?

• Einsteinium-253 is the product of combining 15 neutrons with uranium-238, which then undergoes seven beta decays.
• Einsteinium-253 has a half-life of 20.5 days, according to the Royal Society of Chemistry.
• Scientists at Lawrence Berkeley National Laboratory, the Argonne National Laboratory, and the Los Alamos Scientific Laboratory published the discovery of einsteinium and fermium on June 9, 1955, for the United States Atomic Energy Commission.
• Einsteinium-252 is the most stable isotope of einsteinium and has a half-life of about 471.7 days, according to the Jefferson Laboratory.
• According to the Royal Society of Chemistry, einsteinium has no uses other than scientific research.
• Einsteinium is an actinide element, according to Lenntech, and is found on the bottom row of the periodic table. Actinide elements are attacked by oxygen, steam and acids, but not by alkali metals, such as lithium, sodium, potassium, rubidium, cesium, and francium.
• According to the Los Alamos National Laboratory, einsteinium was the seventh transuranic element to be discovered. Transuranic elements are artificially made, radioactive elements, according to the U.S. Nuclear Regulatory Commission.
• Enough einsteinium had been collected to be visible to the naked eye by 1961 by researchers in Berkeley, California, according to the Royal Society of Chemistry. The amount weighed approximately ten millionths of a gram (1.0 x 10-5 grams or 3.5 x 10-7 ounces).
• Einsteinium is created in very small amounts from bombarding plutonium with neutrons in a nuclear reactor, according to the Royal Society of Chemistry.
• Einsteinium is soft and silver in color, according to Elements Database.
• Einsteinium glows blue in the dark due to the great release of energy as it undergoes radioactive decay, according to Redfern.
• Einsteinium is highly radioactive, according to Lenntech, but as it is not a naturally occurring element, there are no known health risks to the general population. Those who work closely with einsteinium in a laboratory, however, must take precautions to protect themselves from the radiation.
• Due to the rapid decay of einsteinium, it is difficult to study the pure element, according to Redfern. Einsteinium decays into berkelium and californium, causing the nearly all samples of einsteinium to be contaminated.
• Some early reports about the discovery of element 99 unofficially named the element Athenium, for the Greek capital of Athens, according to van der Krogt. The element was officially named for Albert Einstein.
• Einstein was an ironic choice for the name of einsteinium, according to Redfern. Einstein was a pacifist and opposed to the hydrogen bomb.

Current research

There is very little research involving einsteinium. According to Redfern, the main use of einsteinium is to create heavier elements, including mendelevium. Due to the high rate of decay and radioactive nature, there are currently no other uses for einsteinium.

Additional resources

• Peter van der Krogt’s Elementymology website
• Los Alamos National Laboratory
• Images of Elements: Einsteinium

Plutonium is a radioactive, silver metal that can be used to create or destroy. While it was used for destruction soon after it was made, today the element is used mostly for creating energy throughout the world. 

Plutonium was first produced and isolated in 1940 and was used to make the "Fat Man" atomic bomb that was dropped on Nagasaki at the end of World War II, just five years after it was first produced, said Amanda Simson, an assistant professor of chemical engineering at the University of New Haven. 

Just the facts 

Here are the properties of plutonium, according to the Los Alamos National Laboratory:

• Atomic number: 94
• Atomic symbol: Pu
• Atomic weight: 244
• Melting point: 1,184 F (640 C)
• Boiling point: 5,842 F (3,228 C)

Discovery & history

Plutonium was discovered in 1941 by scientists Joseph W. Kennedy, Glenn T. Seaborg, Edward M. McMillan and Arthur C. Wohl at the University of California, Berkley. The discovery occurred when the team bombarded uranium-238 with deuterons that had been accelerated in a cyclotron device, which created neptunium-238 and two free neutrons. The neptunium-238 then decayed into plutonium-238 through beta decay. 

This experiment wasn't shared with the rest of the scientific community until 1946, after World War II. Seaborg submitted a paper on their discovery to the journal Physical Review in March 1941, but the paper was removed when it was discovered that an isotope of plutonium, Pu-239, may be used to create an atomic bomb. 

Soon, Seaborg was sent to lead the Plutonium Production Lab, also known as the Met Lab, at the University of Chicago, according to the Los Alamos National Laboratory. The purpose of the lab was to create plutonium as part of the Manhattan Project. The Manhattan Project was a secret venture during World War II that worked exclusively to develop an atomic bomb. 

On August 18, 1942, they had their first big success. They were able to create a trace quantity of plutonium that was visible to the eye. It equaled only around 1 microgram. From the tiny sample, scientist determined plutonium's atomic weight. 

The Manhattan Project eventually produced enough plutonium for the "Trinity Test." During the test, the world's first atomic bomb, or the "The Gadget," was exploded near Socorro, New Mexico, on July 16, 1945, by Los Alamos Laboratory director Robert Oppenheimer and Army General Leslie Groves. 

Of the test, Oppenheimer said, "We knew the world would not be the same. A few people laughed, a few people cried. Most people were silent. I remembered the line from the Hindu scripture, the Bhagavad-Gita. Vishnu is trying to persuade the Prince that he should do his duty and to impress him takes on his multi-armed form and says, 'Now I am become Death, the destroyer of worlds.' I suppose we all thought that, one way or another," according to the Royal Society of Chemistry. 

The explosion had the energy equivalent of approximately 20,000 tons of TNT. The first war-use atomic bomb dropped on Hiroshima, Japan, on August 6, 1945. That atomic bomb, dubbed "Little Boy," had a uranium core, though. The second bomb, dropped on Nagasaki, Japan, in August 9, 1945, had a plutonium core. The "Fat Man," as it was called, hastened the end of World War II.

Properties of plutonium

Freshly prepared plutonium metal has a silvery bright color but takes on a dull gray, yellow, or olive green tarnish when oxidized in air. The metal quickly dissolves in concentrated mineral acids. A large piece of plutonium feels warm to the touch because of the energy given off by alpha decay; larger pieces can produce enough heat to boil water. At room temperature alpha-form plutonium (the most common form) is as hard and brittle as cast iron. It can be alloyed with other metals to form the room-temperature stabilized delta form, which is soft and ductile. Unlike most metals, plutonium is not a good conductor of heat or electricity. It has a low melting point and an unusually high boiling point. 

Plutonium can form alloys and intermediate compounds with most other metals, and compounds with a variety of other elements. Some alloys have superconductive abilities and others are used to make nuclear fuel pellets. Its compounds come in a variety of colors, depending on the oxidation state and how complex various ligands are. In aqueous solution there are five valance ionic states.

Plutonium, along with all of the other transuranium elements, is a radiological hazard and must be handled with specialized equipment and precautions. Animal studies have found that a few milligrams of plutonium per kilogram of tissue are lethal.

Sources

Plutonium generally isn't found in nature. Trace elements of plutonium are found in naturally occurring uranium ores. Here, it is formed in a way similar to neptunium: by irradiation of natural uranium with neutrons followed by beta decay.

Primarily, however, plutonium is a byproduct of the nuclear power industry. Each year, around 20 tons of plutonium is produced, according to the Los Alamos National Laboratory. Spent nuclear fuel can also be reprocessed to separate usable plutonium from other elements in the fuel.

Atmospheric weapons testing in the 1950s and 1960s left tons of plutonium in the Earth's atmosphere that is still there today, according to the World Nuclear Association.

Uses

For the most part, plutonium isn't used for much. In fact, of the five common isotopes, only two of plutonium's isotopes, plutonium-238 and plutonium-239, are used for anything at all. 

Plutonium-238 is used to make electricity for space probes using radioisotope thermoelectric generators. These generators are switched on when the probes can't get enough solar power because they have traveled too far away from the sun. Some probes that use plutonium-238 are Cassini and Galileo. 

When concentrated enough, plutonium-239 undergoes a fission chain reaction. Because of this, it is used in nuclear weapons and some nuclear reactors.

In fact, one of the biggest uses for plutonium is energy. According to the World Nuclear Association, over one-third of the energy produced in most nuclear power plants comes from plutonium. Plutonium is the main fuel in fast neutron reactors.

Who knew?

For decades, scientists wondered why plutonium didn't act like other metals in its group. For instance, plutonium is a poor conductor of electricity and it doesn't stick to magnets. Now researchers have figured out where its "missing magnetism" has been hiding out and it has to do with the wacky behavior of the electrons in the element's outer shell. Unlike other metals, which have a set number of electrons in their outer shells, when in a ground state, plutonium can have four, five or six electrons there.

This fluctuating number of outer-shell electrons explains why plutonium isn't magnetic: In order for an atom to interact with magnets the unpaired electrons in its outer shell must line up in a magnetic field. [Read more about plutonium's missing magnetism]

Plutonium's most stable isotope, plutonium-244, can last a long time. It has a half-life of about 82 million years and decays into uranium-240 through alpha decay, according to the Jefferson Lab.

Plutonium was named after the planet, Pluto. This is because it came after Uranium, which was named after the planet Uranus, and neptunium, which was named after the planet Neptune.

Plutonium emits neutrons, beta particles and gamma rays.

Additional resources

• U.S. Nuclear Regulatory Commission: Backgrounder on Plutonium
• NASA: What is Plutonium-238?
• Centers for Disease Control and Prevention: Plutonium Radioisotope Brief

Oganesson is a radioactive, artificially produced element about which little is known. It is expected to be a gas and is classified as a non-metal. It is a member of the noble gas group.

The element, No. 118 on the Periodic Table of Elements, had previously been designated ununoctium, a placeholder name that means one-one-eight in Latin. In November 2016, the International Union of Pure and Applied Chemistry (IUPAC) approved the name oganesson for element 118.

The IUPAC also approved names for elements 113 (nihonium, with atomic symbol Nh), 115 (moscovium, Mc) and 117 (tennessine, Ts).

The name oganesson honors Yuri Oganessian "for his pioneering contributions to transactinide elements research," IUPAC officials said, referring to elements with atomic numbers 104 through 120. "His many achievements include the discovery of super-heavy elements and significant advances in the nuclear physics of super-heavy nuclei, including experimental evidence for the 'island of stability,'" an idea suggesting that super-heavy elements can become stable at some point in their existence.

Just the facts

Atomic Number: 118 Atomic Symbol: Og Atomic Weight: [294] Melting Point: Unknown Boiling Point: Unknown

Discovery

Oganesson was discovered in 2002 by Russian scientists at the Joint Institute for Nuclear Research in Dubna, Russia. Three years prior, in 1999, the team at the Lawrence Berkeley Labs in California published a paper announcing the discovery of element 118, but their results could not be replicated and the team retracted their paper. In 2006, the element was officially announced by the Dubna team and by the Lawrence Livermore National Laboratory team, who had been working with the Dubna scientists.

Properties of oganesson

Oganesson has one known isotope, ²⁹⁴Og, with a half-life of about 0.89 milliseconds. Through alpha decay, it turns into ²⁹⁰Lv (livermorium-290).

The atomic weight for manmade transuranium elements is based on the longest-lived isotope. These atomic weights should be considered provisional since a new isotope with a longer half-life could be produced in the future. [See Periodic Table of the Elements]

Sources of oganesson

The Russian scientists who produced oganesson bombarded atoms of californium with ions of calcium for 1,080 hours. This resulted in three atoms of oganesson.

Uses of oganesson

Since only a few atoms of oganesson have ever been made, it has no practical uses outside of scientific study.

Additional resources

• Los Alamos National Laboratory
• Jefferson Lab

Tennessine is a radioactive, artificially produced element about which little is known. It is expected to be a solid, but its classification is unknown. It is a member of the halogen group.

The element, No. 117 on the Periodic Table of Elements, had previously been designated ununseptium, a placeholder name that means one-one-seven in Latin. In November 2016, the International Union of Pure and Applied Chemistry (IUPAC) approved the name tennessine for element 117.

The IUPAC also approved names for elements 113 (nihonium, with atomic symbol Nh), 115 (moscovium, Mc) and 118 (oganesson, Og).

Names for elements 115 and 117 were proposed by their discoverers at the Joint Institute for Nuclear Research in Dubna, Russia; the Oak Ridge National Laboratory in Tennessee; Vanderbilt University in Tennessee; and Lawrence Livermore National Laboratory in California. Both element names, moscovium and tennessine, honor regions where experiments linked to creating the elements took place.

Just the facts

Atomic Number: 117 Atomic Symbol: Ts Atomic Weight: [294] Melting Point: Unknown Boiling Point: Unknown

Discovery

Element 117 was discovered in 2010 and jointly announced on April 5 of that year by scientists at the Joint Institute for Nuclear Research in Dubna, Russia, and scientists at the Lawrence Livermore National Laboratory in California.

Properties of tennessine

Tennessine has two isotopes with known half-lives and two with unknown half-lives. The most stable isotope is ²⁹⁴Ts, with a half-life of about 80 milliseconds.  It decays through alpha decay. Tennessine’s other isotopes are suspected to decay through both alpha decay and spontaneous fission.

The atomic weight for manmade transuranium elements is based on the longest-lived isotope. These atomic weights should be considered provisional since a new isotope with a longer half-life could be produced in the future. [See Periodic Table of the Elements]

Sources of tennessine

The scientists who created tennessine bombarded atoms of berkelium with ions of calcium until atoms tennessine were produced.

Uses of tennessine

Since only a few atoms of tennessine have ever been made, it has no practical uses outside of scientific study.

Additional resources

• Los Alamos National Laboratory
• Jefferson Lab

Moscovium is a radioactive, synthetic element about which little is known. It is classified as a metal and is expected to be solid at room temperature. It decays quickly into other elements, including nihonium.

The element had previously been designated ununpentium, a placeholder name that means one-one-five in Latin. In November 2016, the International Union of Pure and Applied Chemistry (IUPAC) approved the name moscovium for element 115.

The IUPAC also approved names for elements 113 (nihonium, with atomic symbol Nh), 117 (tennessine, Ts) and 118 (oganesson, Og).

Names for elements 115 and 117 were proposed by their discoverers at the Joint Institute for Nuclear Research in Dubna, Russia; the Oak Ridge National Laboratory in Tennessee; Vanderbilt University in Tennessee; and Lawrence Livermore National Laboratory in California. Both element names, moscovium and tennessine, honor regions where experiments linked to creating the elements took place.

Just the facts

Atomic Number: 115 Atomic Symbol: Mc Atomic Weight: [288] Melting Point: Unknown Boiling Point: Unknown

Discovery

Moscovium was discovered in 2003 and officially announced on Feb. 2, 2004. It was created and announced by scientists at the Joint Institute for Nuclear Research in Dubna, Russia, and scientists at the Lawrence Livermore National Laboratory in the United States.

Properties

Moscovium has four isotopes with known half-lives, the most stable of which is ²⁸⁹Mc, with a half-live of about 220 milliseconds.

The atomic weight for manmade transuranium elements is based on the longest-lived isotope. These atomic weights should be considered provisional since a new isotope with a longer half-life could be produced in the future. [See Periodic Table of the Elements]

Sources of moscovium

To make moscovium, the scientists in Russia and the United States bombarded atoms of americium with ions of calcium in a cyclotron. This produced four atoms of moscovium.

Uses of moscovium

Only a few atoms of moscovium have ever been made, and they are only used in scientific study. It is used to make nihonium.

Additional resources

• Los Alamos National Laboratory
• Jefferson Lab

Nihonium is a radioactive, synthetic element about which little is known. It is classified as a metal and is expected to be solid at room temperature.

The element, No. 113 on the Periodic Table of Elements, had previously been designated ununtrium, a placeholder name that means one-one-three in Latin. In November 2016, the International Union of Pure and Applied Chemistry (IUPAC) approved the name nihonium for element 113.

Scientists with Japan's RIKEN Nishina Center for Accelerator-Based Science proposed the element name nihonium, which is one way to say "Japan" in Japanese and means "the land of the rising sun," according to the IUPAC. Nihonium's atomic symbol is Nh.

The IUPAC also approved names for elements  115 (moscovium, with atomic symbol Mc), 117 (tennessine, Ts) and 118 (oganesson, Og).

Discovery

Kosuke Morita and his colleagues created the elusive element on Aug. 12, 2012, after colliding zinc nuclei together in a thin layer of bismuth. Like other superheavy elements, after 113 was created, it quickly decayed, ultimately turning element 113 into 111, and then 109, 107, 105, 103 and finally into element 101, according to Morita.

Nihonium has six isotopes with known half-lives. The most stable is ²⁸⁶Nh, with a half-life of about 20 seconds.

The atomic weight for manmade transuranium elements is based on the longest-lived isotope. These atomic weights should be considered provisional since a new isotope with a longer half-life could be produced in the future. [See Periodic Table of the Elements]

Just the facts

Atomic Number: 113 Atomic Symbol: Nh Atomic Weight: [286] Melting Point: Unknown Boiling Point: Unknown

Sources

To produce nihonium, Element 115 — now called moscovium — must first be made. Atoms of americium (Element 95) are bombarded with ions of calcium (Element 20) in a cyclotron, which produces moscovium. Then, moscovium quickly alpha decays into nihonium.

Uses of nihonium

Only a few atoms of nihonium have ever been made, and they are only used in scientific study.

Additional resources

• Los Alamos National Laboratory
• Jefferson Lab

Iridium is the most corrosion-resistant element on the Periodic Table of Elements. It also has the highest density of all the elements. Because it resists corrosion, it is used to set standards in weights and measures. But because it is so dense and brittle, it is hard to machine, form or work it unless it is heated to extreme temperatures.

Properties

Iridium is a member of the platinum family and is white in color with a yellowish hue. It has a density of 22.65 grams per cubic centimeter. By comparison, the density of lead is 11.34 g/cm³ and the density of iron is 7.874 g/cm³. 

Iridium is not affected by acids, bases, or most other strong chemicals, according to Chemistry Explained. That property makes it useful in making objects that are exposed to such materials.

Just the facts 

Here are other properties of iridium, according to the Los Alamos National Laboratory:

• Atomic number (number of protons in its nucleus): 77
• Atomic symbol (on the Periodic Table of Elements): Ir
• Atomic weight (average mass of the atom): 192.217
• Melting point: 4,435 F (2,446 C)
• Boiling point: 8,002.4 F (4,428 C)
• Stable isotopes: 2, which are iridium-191 (37.3 percent) and iridium-193 (62.7 percent)

History

Several chemists may have discovered iridium about the same time in 1803, according to an article in the journal Platinum Metals Review. English chemist Smithson Tennant, French chemists H.V. Collet-Descotils, A.F. Fourcroy and N.L. Vauquelin all are said to have found iridium in the acid-insoluble residues of platinum ores. Tennant usually gets the credit, though.

Tennant discovered iridium by dissolving crude platinum in diluted aqua regia (a mixture of nitric and hydrochloric acids), then by treating the black residue left behind in turn with alkalis and acids, according to the Royal Society of Chemistry. After this treatment, the residue separated into two new elements. At the Royal Institution in London he announced his findings and named one element iridium and the other osmium. The name iridium comes from the Latin word iris, which means rainbow. Though the metal itself isn't rainbow colored, it is called this because of its multi-colored compounds.

Because iridium is very resistant to corrosion, the standard meter bar was made of 90 percent platinum and 10 percent iridium. This bar was replaced as the definition of a meter in 1960, though. The meter was redefined in terms of the orange-red spectral line of krypton. However, the international prototype kilogram, which defines a kilogram, also made of a platinum and platinum-iridium alloy, is still in use around the world.

Sources

Today, iridium is commercially recovered as a byproduct of copper or nickel mining. Ore containing iridium is found in Brazil, the United States, Myanmar, South Africa, Russia and Australia. 

Pure iridium is so rare on the Earth's crust that there is only about 2 parts per billion located in the crust, according to Chemistry Explained. 

"Iridium is one of the densest and rarest of Earth's natural elements. It is so dense that it mainly exists in the Earth's core, rather than crust," said Amanda Simson, an assistant professor of chemical engineering at the University of New Haven.

But some iridium exists in the crust. In 1980 scientist Luis Alvarez and his son Water Alvarez found significant amounts of iridium in a certain part of the Earth's crust, spread out all over the Earth's surface. "They speculated that it was caused by a meteor and linked this to the extinction of dinosaurs 66 million years prior," explained Simson. 

Uses

Though brittle, iridium can be worked if heated to a white heat of 2,200 to 2,700 degrees Fahrenheit (1,200 to 1,500 degrees Celsius), according to Encyclopedia Britannica. Iridium's principal use is to harden platinum by making a platinum alloy. 

It is also used to make devices needed for high temperatures and in electrical contacts. It is also used on some optical lenses to reduce glare. A compound of osmium and iridium, called osmiridium, is used in fountain pen tips and compass bearings. Super-strong jewelry is also made of an iridium and platinum alloy. 

Additional resources

• Jefferson Lab: The Element Iridium
• US National Library of Medicine: Studies on Some Iridium(III) Complexes with Schiff Bases Derived from Amino Carboxylic Acids
• Iridium Complexes in Organic Synthesis

Tungsten is known as one of the toughest things found in nature. It is super dense and almost impossible to melt. Pure tungsten is a silver-white metal and when made into a fine powder can be combustible and can spontaneously ignite. Natural tungsten contains five stable isotopes and 21 other unstable isotopes.

Tungsten is used in many different ways because it is very strong and durable. It is very resistant to corrosion and has the highest melting point and highest tensile strength of any element. Its strength comes when it is made into compounds, though. Pure tungsten is very soft.

Just the facts

Here are the properties of tungsten, according to the Los Alamos National Laboratory:

• Atomic number: 74
• Atomic symbol: W
• Atomic weight: 183.84
• Melting point: 6,192 F (3,422 C)
• Boiling point: 10,030 F (5,555 C)

History

The first use of tungsten was more than 350 years ago. Chinese porcelain makers used a tungsten pigment that was a unique peach color, according to the Royal Society of Chemistry.

Much later, in 1779, Peter Woulfe examined a mineral from Sweden and realized it contained a new type of metal, but that’s about as far as the research went. In 1781, Wilhelm Scheele continued the research on this new metal and isolated an acidic white oxide. Neither one of these men are credited with the element’s discovery, though.

Juan and Fausto Elhuyar get that honor. At the Seminary at Vergara in Spain, they researched this mysterious metal. In 1783 they isolated the metal oxide from wolframite and then, unlike the others, reduced it to tungsten metal by heating it with carbon.

Sources

Most tungsten resources are found in China, South Korea, Bolivia, Great Britain, Russia and Portugal, as well as in California and Colorado. Though it is found in these many places, 80 percent of world’s supply is controlled by China, according to the BBC. 

The element naturally occurs in the minerals scheelite, wolframite, huebnertie and ferberite. It is harvested from the minerals by reducing tungsten oxide with hydrogen or carbon.

Once it is sourced, tungsten is often mixed into alloys. The hardest alloys are shaped using diamonds. Diamonds are the only things harder than some tungsten alloys.

Uses

One of the most common, and hardest, tungsten compounds is tungsten carbide. Because of its strength when made into compounds, tungsten is used to harden saw blades and make drill bits. It can take around 10 minutes to cut just one drill bit from tungsten using a diamond cutting system, according to the BBC. Some jewelers also use tungsten carbide to make wedding bands and other rings.

Another tungsten compound that is particularly useful is tungsten disulfide. It is used as a dry lubricant in temperatures as high as 932 degrees Fahrenheit (500 degrees Celsius), according to the Jefferson Lab.

Some other uses of tungsten include metal evaporation work, the manufacturing of paints, making glass-to-metal seals and creating electron and television tubes.

The military uses tungsten to make bullets and missiles used in “kinetic bombardment.” This type of attack uses a super dense material to breach armor instead of explosives.

Its resistance to heat is helpful when using it in the heating elements for electrical furnaces, spacecraft applications, welding and other high-temperature applications. It was also used in making different types of lighting for this reason. The hotter a filament can get without melting, the brighter the bulb. In 1908 inventor William D. Coolidge discovered that tungsten was an ideal filament material. Today, though, most bulbs use more energy efficient materials. It is still used in X-ray filaments and in electrical contacts of various electronics, however.

Biologically, some bacteria use tungsten to reduce carboxylic acids to aldehydes.

Who knew?

This element is used for trickery. “Tungsten may not have gold’s luster, but it does have its density (within 0.36 percent) which means that if you cover a brick of tungsten with a coating of gold – and you test the brick to see if it weighs as much as gold – it will be almost correct,” Amanda Simson, an assistant professor of chemical engineering at the University of New Haven, told Live Science. “Thus, tungsten has been found in counterfeit gold bricks.”

Tungsten comes from a Swedish term, tung sten, that means "heavy stone.”

Tungsten's chemical symbol is a W, which may seem weird since there isn’t a W in the word. The W actually comes from the element’s other name, wolfram. The name wolfram comes from the mineral the element was discovered in, wolframite. Wolframite means "the devourer of tin," which is appropriate since the mineral interferes with the smelting of tin.

Additional resources

• Biological Chemistry: Purification and Some Properties of the Tungsten-Containing Carboxylic Acid Reductase from Clostridium formicoaceticum
• Proceedings of the National Academy of Science: Identification and Characterization of the Tungsten-Containing Class of Benzoyl-Coenzyme A Reductase
• Centers for Disease Control and Prevention: Tungsten Toxicity

This article was updated on Feb. 3, 2020 to correct the boiling point of tungsten. 

Radium is a highly radioactive element and can be extremely dangerous. However, it was once used in many everyday products, including wristwatches and toothpaste, and thought to have curative properties until its intense radioactivity was found to cause adverse health effects.

Radium has an abundance of about 1 part per trillion in the Earth's crust, according to Chemicool. Trace amounts of radium are found in uranium ore, because radium is created from the decay of the uranium atom, which then into several other unstable elements before finally ending in the element lead. There are several known isotopes of radium, but due to the rapid decay rates of many of the isotopes, it is uncertain about the natural abundances of the radium isotopes.

Just the facts

• Atomic number (number of protons in the nucleus): 88
• Atomic symbol (on the periodic table of elements): Ra
• Atomic weight (average mass of the atom): 226
• Density: 3.2 ounces per cubic inch (5.5 grams per cubic cm)
• Phase at room temperature: solid
• Melting point: 1,292 degrees Fahrenheit (700 degrees Celsius)
• Boiling point: 2,084 F (1,140 C)
• Number of natural isotopes (atoms of the same element with a different number of neutrons): 33
• Most common isotopes: Ra-226 (unknown percent of natural abundance), Ra-223 (unknown percent of natural abundance), Ra-224 (unknown percent of natural abundance), Ra-228 (unknown percent of natural abundance)

History

Marie and Pierre Curie, Polish and French chemists, discovered radium in 1898, according to New World Encyclopedia. The discovery came from the study of pitchblende (a type of uranium ore) found in Bohemia (today's Czech Republic). The uranium was removed from the ore and the remains were found to still be radioactive. The radioactive remains were then separated out and when the spectrum was studied, the material was found to be primarily barium with an unknown element.

According to Peter van der Krogt, a Dutch historian, the element was named for the Latin word "radius" or "ray" because the radiation emitted from the new element was about 3 million times greater than the radiation from uranium. The Curies were able to extract about 1 milligram of radium from nearly 10 tons of pitchblende, according to the Royal Society of Chemistry.

Pure radium was isolated in 1902 by electrolysis by Marie Curie and Andre Debierne, a French chemist, according to New World Encyclopedia. Radium E, known to be bismuth-210, was the first synthetic radioactive element that was created synthetically by scientists at the University of California, according to Time.

Who knew?

• According to Chemicool, radium has an abundance in the Earth's crust about 1 part per trillion by weight. This makes it the 84thmost abundant element in Earth's crust according to Periodic Table.
• Radium is the heaviest alkaline earth metal, according to Encyclopedia. The other alkali earth metals include beryllium, magnesium, calcium, strontium and barium.
• Radium changes from a silvery white color to black when exposed to air, according to Lenntech due to oxidation.
• According to Chemicool, the radium isotope that has the longest half-life is radium-226 with a half-life of 1602 years.
• According to the Agency for Toxic Substances and Disease Registry, radium typically enters the body when it is breathed in or swallowed. Health effects from radium exposure include cancer, anemia, cataracts, and death.
• Radium emits alpha particles (two protons and two neutrons bonded together), beta particles (high energy electrons or positrons), and gamma rays (the most energetic wavelength of light), according to New World Encyclopedia.
• According to the Royal Society of Chemistry, radium is in the same group as calcium and is sometimes used to target bone cancer. Alpha particles emitted by the radium kills the cancerous cells.
• Radium is primarily extracted as a byproduct in uranium mining, according to the Royal Society of Chemistry. Most of the radium comes from uranium mines in Democratic Republic of Congo and Canada.
• According to Chemistry Explained, radium is extracted today from uranium ores in much the same way that Marie and Pierre Curie did in the late 1890s and early 1900s.
• According to Periodic Table, radium was used in clocks to paint the numbers and the hands so that they were visible in the dark. This practiced ceased after too many factory workers died from exposure.
• According to Encyclopedia, radium combines with nearly all non-metals including oxygen, fluorine, chlorine, and nitrogen.
• A curie (Ci) is a unit that is named for the amount of radiation emitted of a quantity of radionuclides that is equal to one gram of radium, or at a rate of decays at 37 billion disintegrations per second, according to the U.S. Nuclear Regulatory Commission.
• According to Lenntech, because radium is naturally occurring in small amounts in our environment, we are constantly exposed to small amounts of radiation. There is no evidence that the radium radiation levels are harmful.
• Marie and Pierre Curie's laboratory notebooks are still too radioactive due to their work with radium to be handled today, according to the Jefferson Lab.
• Due to her work with radium, Marie Curie was the first woman to win a Nobel Prize in physics (1903) and the first scientist to win two Nobel Prizes (second in 1911), according to Biography.
• According to the Los Alamos National Laboratory, radium is used to produce radon gas, which is typically used to treat several diseases including cancer.
• Radium is an unstable element and undergoes several stages of radioactive decay reaching its end product of lead, according to New World Encyclopedia. 

Current research

Radium is often used to treat various forms of cancers. In a study published in the Journal of Nuclear Medicine, Isis Gayed and other researchers from Texas discuss treating bony metastases brought on by prostate cancer. The patients were treated with isotope Ra-223 and closely followed through treatments. Several factors were considered and compared before and after treatments including pain levels, PSA, creatinine, and hematological values. The patients that had lower PSA and creatinine levels and higher hemoglobin levels responded better to the radium treatments.

A 2014 article by Ashley Lehman, an American researcher, published in the Journal of Nuclear Medicine, discusses how the medication used in the above study, radium-223 dichloride (Xofigo), works. When radium encounters bone, it behaves similarly to calcium and gravitates toward where new bone formation is occurring. The radium in the medication collects in the bony metastatic sites there the damage from alpha particles being emitted by the decaying radium is primarily limited to the surrounding cancerous tissue as the alpha particles only travel short distances. The study concluded that using radium-223 was a promising treatment for those with prostate cancer and that the same treatment was currently being studied with in breast cancer patients who have developed bone metastases.

There are several ongoing trials for using radium-223 for breast cancer, including a double-blind, placebo-controlled trial by a group of American researchers, with an abstract published in the American Association for Cancer Research in 2015. The study is aiming to get a total of 227 participants to study the effects and safety of using radium to treat bony metastatic sites caused by breast cancer.

Additional resources

• Lenntech
• Los Alamos National Laboratory
• Chemistry Explained

Lead is an incredibly useful metal, but it is also toxic to humans. In fact, if we didn't have to worry about breathing in its dust or ingesting its particles, lead would be in widespread use due to its highly industry-friendly properties, such as excellent malleability and corrosion resistance.

Throughout history — before the scientific advancements of the 20th century revealed its potent toxicity — lead was widely used in a variety of products, including cosmetics, paint, solder, pipes and gasoline. Certain properties of lead, namely its ductility and resistance to corrosion and tiny leaks, make it a particularly good material for constructing water pipes. Even the ancient Romans made their water pipes out of lead, causing some to believe that lead poisoning, at least partially, led to the fall of the Roman Empire. 

Natural element

Lead is a highly lustrous, bluish-white element that makes up only about 0.0013 percent of the Earth's crust, according to the Jefferson Lab. It is not considered rare, however, since it is fairly widespread and easy to extract. Lead typically occurs in very small amounts in ores such as galena, anglesite and cerussite. Lead is commonly mined and smelted in Missouri, Idaho, Utah, Colorado, Montana and Texas, according to Plumbing Manufacturers International. About one-third of the lead in the United States is recycled.

The chemical symbol for lead is Pb, which comes from the Latin word plumbum, meaning "waterworks," referring back to ancient times when the metal was widely used in the construction of water pipes. Although there are 27 lead isotopes, only four are considered stable.

Lead linings

Although lead has been phased out of many of its previous uses, this non-corrosive metal is actually quite useful in products that hold or touch highly acidic substances. For example, lead is used to line tanks that hold corrosive liquids, such as sulfuric acid. It is also used in lead-acid storage batteries, such as those found in automobiles.

Because of its density and ability to absorb vibration, lead also makes an excellent shield against different types of harmful radiation, such as those found in X-ray machines and nuclear reactors, according to Jefferson Lab. Lead is also still used in some bullets and ammunition. 

Leaded gasoline

Tetraethyl lead was added to gasoline in the 1920s to help reduce engine knocking, wear and tear and pre-ignition. Almost immediately, industry workers started to become extremely ill and some even died. At Dupont's manufacturing plant in New Jersey, it was particularly bad — eight workers died between 1923 and 1925. Finally, after 44 workers at Standard Oil's plant had been hospitalized, public awareness and outcry finally began to gather steam, according to Chemistry LibreTexts. Although the U.S. Public Health Service held a conference in 1925, lead was ultimately allowed to remain in gasoline for decades in spite of all the damage it was causing. It wasn't until the late 1970s that leaded gasoline started to get phased out. It was finally banned for all on-road vehicles in 1996.

Just the facts

• Atomic number (number of protons in the nucleus): 82
• Atomic symbol (on the periodic table of the elements): Pb
• Atomic weight (average mass of the atom): 207.2
• Density: 11.342 grams per cubic centimeter
• Phase at room temperature: Solid
• Melting point: 621.4 degrees Fahrenheit (327.46 degrees Celsius)
• Boiling point: 3,180.2 degrees Fahrenheit (1,749 degrees Celsius)
• Number of isotopes (atoms of the same element with a different number of neutrons): 27; 4 stable 
• Most common isotopes: Pb-208 (52.4 percent of natural abundance); Pb-206 (24.1 percent of natural abundance); Pb-207 (22.1 percent of natural abundance); Pb-204 (1.4 percent of natural abundance)

Toxicity today

Since lead was used in so many products before people recognized the extent of its toxicity, it continues to pose a real public health danger today. For example, children living in older homes can still breathe in or ingest dust or paint chips from peeling lead paint on the walls. Lead can also leach into the drinking water supply through older, corroded lead pipes, faucets and solder. Nearly all homes built before the 1980s have lead solder connecting to copper pipes, according to Plumbing Manufacturers International. Even some major U.S. cities still use lead pipes to carry water from the utilities to our homes and businesses. When the water chemistry is very carefully controlled, however, it keeps lead from leaching into the drinking water.

Most cases of lead poisoning are due to chronic low-dose exposure. Since symptoms of slow lead poisoning are mainly emotional and mental in nature, lead poisoning may be the last thing people suspect. Children are at the greatest risk. Lead can delay physical and mental development in babies and young children. In adults, slow accumulation of lead can result in kidney and nervous system damage, anemia, stroke or cancer, according to Utah State University. 

Quan Lu is an associate professor of environmental genetics and pathophysiology at Harvard T.H. Chan School of Public Health. He is the senior author of a recent study investigating the harmful effects of lead on neural stem cells and children's neurodevelopment. 

"While acute poisoning of lead (Pb) can cause abdominal pain, weight loss, vomiting and even death, the majority cases of Pb exposure in children are chronic and low-dose," said Lu. "Pb exposure in children has been consistently linked to impaired neurological development and cognitive dysfunction as well as persistent antisocial and delinquent behavior." 

"Pb disturbs neuronal function. Pb neurotoxicity is determined by intricate interplays between the metal and target neural cells, and there is overwhelming evidence documenting the detrimental effects of Pb in neurons," he told Live Science.

Lu added that previous studies had shown that lead potently inhibits the N-methyl-D-aspartate receptor (NMDA) receptor, which plays an essential role in brain development, synaptic plasticity and learning & memory. 

"We need to reduce the chances of Pb exposure," he said. "That means we need to eliminate the sources of Pb contamination: in water, in gasoline, in paints and in any consumer products. Recent incidences of Pb contamination in drinking water in several U.S. cities highlight the urgent need to reduce Pb contamination to protect children." 

Who knew?

• Pencils do not — and never did — contain lead. The "lead" in pencils is actually graphite. When large deposits of graphite were first discovered in England in the 1500s, people assumed they had discovered lead, hence the name.
• In ancient alchemy, lead was considered the prima matera, or primal matter. Alchemists associated lead with Saturn, the Roman god of agriculture and harvest. Accordingly, lead was represented with Saturn's symbol, the scythe, ♄.
• In Ancient Rome, lead was added to wine by unscrupulous individuals in order to increase the perception of the wine's sweetness. Although this practice was strictly forbidden in 1498 by the Roman Catholic Church, it continued on for quite some time, resulting in large-scale poisonings well into the late 18th century.
• In 16th and 17th century Europe, lead was used in cosmetics as a way to obtain the fresh "white-faced" look that was so popular at the time, particularly among the aristocracy. One famous user of lead makeup was Queen Elizabeth I of England — her white lead paint was said to give her the coveted "Mask of Youth." Unfortunately, long-term contact with lead often resulted in rotten teeth, and ironically, tooth replacements were often made of lead as well, causing even worse damage to the already lead-poisoned person.

Flint, Michigan: Drinking water crisis

One of the most notorious cases of lead leaching into drinking water occurred recently in Flint, Michigan. In an effort to save money, officials had decided to switch the source of the city's drinking water from the Detroit Water and Sewerage Department (DWSD) to the Karegnondi Water Authority (KWA). In the meantime, however, they would need to pull water from the Flint River, beginning on April 25, 2014.

Within weeks, Flint residents began to complain about the smell and color of their tap water. Tests revealed high levels of E.coli and total coliform bacteria in the water supply, which prompted the city to chlorinate the water at higher-than-usual levels. This chlorination, in addition to the fact that they had not implemented any corrosion protection, caused massive pipe corrosion, allowing lead to leach into the drinking water. 

In many homes, the levels of lead in the drinking water were far above the Environmental Protection Agency's maximum safety level of 15 parts per billion (ppb). In fact, the water in one home was tested by Virginia Tech researchers as having lead levels at 13,200 ppb — over three times the level considered to be hazardous waste. Unfortunately, a child living in that home was diagnosed with lead poisoning.

Additional resources

• Los Alamos National Lab: Lead
• Royal Society of Chemistry: Lead
• Jefferson Lab: The Element Lead