Astronauts have grown a common crystal in space — but this extraterrestrial version is exhibiting a very uncommon behavior. Under a microscope, the crystal looks more like a tiny futuristic sculpture than something that formed naturally. The pyramid-like assembly curls in on itself in stair-stepped layers, seeming almost too orderly to be real. But according to scientists, the striking shape is what happens when gravity stops calling the shots.

The image, shared on July 18 by NASA astronaut Don Pettit, and accompanying video were captured aboard the International Space Station (ISS) and show potassium chloride crystals growing in microgravity — the near-weightless environment experienced in orbit.

Microgravity isn't quite zero gravity; astronauts experience a gravitational pull roughly a million times weaker than on Earth's surface, Anne Wilson, a professor of chemistry and biochemistry at Butler University, told Live Science. That reduction in gravity's tug has profound impacts on the human body, and, as the new image shows, completely changes how liquids and crystals behave as well.

"[Other] forces start to become far more pronounced than gravity forces," Wilson said, pointing to molecular attraction and polarity as two influences that suddenly take over once gravity fades into the background.

Those conditions enable potassium chloride — a common salt used as a sodium substitute in foods and sports drinks — to crystallize in ways that are nearly impossible to reproduce in a terrestrial lab.

On Earth, growing crystals eventually get heavy enough to sag, break or settle at the bottom of their container. But in microgravity, they keep growing outward from wherever they started.

Potassium chloride naturally forms cube-shaped crystals, Wilson said. In microgravity, those crystals grow mainly from their edges and corners instead of across their faces, leaving the center hollow and creating the striking pyramid-like structures seen in the image.

Scientists call this pattern hopper growth, and it isn't unique to potassium chloride (ordinary table salt forms the same way). As new corners emerge during growth, the crystal can change direction entirely, producing the elegant scrolling pattern visible in the new image.

While the crystal is visually stunning, the science behind it could prove just as valuable. Materials grown without gravity's interference often form with fewer defects, giving researchers a rare look at what a truly "ideal" structure looks like. This insight could help refine manufacturing processes for semiconductors and other advanced materials back on Earth.

For Wilson, images like this serve another purpose: inspiring curiosity.

"I love seeing videos of how things behave in space," she said. "Who would think that something as simple as potassium chloride could still be super cool?"

ISS astronauts spend a significant amount of their time conducting science experiments that would be near-impossible on Earth. Recently, astronauts used the space station’s newly upgraded Cold Atom Laboratory to create and study an elusive fifth state of matter called the Bose-Einstein condensate. And last year, researchers demonstrated how moss spores exposed to space for nine months on the outside of the ISS continued to grow after being returned to Earth.

A mysterious wavelength of light is missing from the dwarf planet Pluto and Saturn's supersized moon Titan, new James Webb Space Telescope (JWST) observations show. The surprising discovery hints that these worlds harbor an unknown molecule that has not been seen in any other solar system world or exoplanet so far.

Every element or molecule in the universe absorbs unique wavelengths of electromagnetic radiation. Therefore, one of the main ways astronomers study distant worlds — both inside and outside the solar system — is by closely examining the light that reflects off them and searching for dark "absorption lines" that correspond to the wavelengths of known chemical compounds.

For example, molecular oxygen absorbs light at 230 nanometers, so if the electromagnetic spectrum of a faraway exoplanet has an absorption line at this frequency, researchers can be confident that its atmosphere contains oxygen, according to a 2021 study.

JWST has proved to be exceptionally good at capturing electromagnetic spectra and identifying specific chemicals in exoplanet atmospheres, around distant stars and within primitive galaxies. It has even identified a molecule on one alien world that could point to extraterrestrial life.

In a new study, uploaded June 11 to the preprint server arXiv, researchers analyzed JWST data from Pluto and Titan, focusing on very small wavelengths that have been relatively unexplored until now. This revealed a specific absorption line at around 5.11 micrometers in both worlds' spectra. (These findings have not been published in a peer-reviewed journal yet.)

The team pored through similar studies on other planetary spectra but "did not find any band referenced in these publications that corresponds to the location of the observed absorption in Titan and Pluto," the researchers wrote in the paper.

A molecular mystery

The discovery is even stranger because there are very few similarities between Pluto and Titan that could explain why they share a molecule not found anywhere else.

Titan is the largest of Saturn's many moons and is even larger than Mercury. It is also the only solar system world, other than Earth, that is known to have liquid rivers and oceans on its surface. Pluto, on the other hand, is a completely frozen world that's around half the size of Titan and roughly four times farther from the sun than Saturn's satellite is.

Both worlds do have similar atmospheres that are rich in methane and nitrogen. However, the researchers are confident that the molecule responsible for the 5.11-micrometer absorption line is located on both worlds' surfaces, not in their atmospheres.

Pluto's absorption line is around three times thicker than Titan's, meaning the mystery molecule is likely much more abundant on the dwarf planet. But on Titan, the molecule seems to be unevenly distributed, with a stronger absorption line on its trailing side — the hemisphere opposite to its forward momentum around Saturn — than on its leading side.

The researchers proposed that it could be benzene — a ring-shaped hydrocarbon — mixed with an unknown molecule, or some form of acetylene or ketene ice. However, much more work is needed to prove that any of these potential candidates are responsible for absorbing this specific wavelength, they wrote.

NASA's Dragonfly spacecraft, which is set to launch no earlier than 2028 and fly through Titan's atmosphere in 2034, could eventually shed more light on the situation. The helicopter-like craft's onboard spectrograph could identify the mystery molecule on Saturn's moon, which would also help reveal if it is viable on Pluto, the researchers suggested. But in the meantime, we'll have to wait to unravel this curious cosmic conundrum.

See how well you know our planetary neighborhood with our solar system quiz!

If you push two metal plates together on Earth, nothing happens. But if you take those same plates into the vacuum of space, they can fuse into a single piece of metal.

This phenomenon, called cold welding, has been a known hazard for spacecraft engineers for a long time. So what's actually happening at the atomic level, and why does space make it so much easier?

The answer comes down to a lack of oxygen in space, experts told Live Science.

Metals are made of lattices — structures where atoms are bonded to one another. But atoms near the surface of a metal aren't bonded to anything on the outward-facing side. If given the chance, they'll "reach out" and share electrons with the surface of another piece of metal.

But on Earth, nearly every metal surface is coated in an oxide layer. The layer is just a few atoms thick and forms when metal meets oxygen. "Once the oxide is formed, it's over," Julia Greer, a materials scientist at Caltech, told Live Science. "Then it can't cold weld anymore, because the oxygen basically passivates these bonds."

That thin oxide acts like an insulating wrapper. Without it, the free electrons at the surface of one metal piece stop recognizing which atom they belong to. "Those electrons don't know if it's in this piece or if it's in that piece, so they begin sharing the electrons, and essentially that cold welds things together," Sven Bilén, a professor of engineering design and aerospace engineering at Penn State, told Live Science.

In space, there's no oxygen to rebuild that layer once it's gone. The cold and the radiation make things worse. Bombardment from solar and ionic radiation in orbit can scour metal surfaces clean, Greer said, leaving freshly exposed atoms primed to bond. "Everything in space is conducive to cold welding," she said.

Metal surfaces are never perfectly smooth, either. At a microscopic level, they're jagged ‪—‬ more like tiny mountain ranges than flat plains, said Zachary Cordero, an aerospace engineer at MIT.

Pressing two surfaces together, especially with any sliding or vibration, can shear off the oxide layer that formed while the metal was on Earth and flatten those peaks into metal-to-metal contact. "You're breaking up the surface oxide, and you're forming metallurgical bonds," Cordero said.

Why cold welding worried early spacecraft engineers

Cold welding in space has long been a problem. "If there is cold welding, things can become stuck in place," Cordero told Live Science. "If you have a deployable structure and there's cold welding, you might freeze the mechanism, or a door might become locked, or something might become immobilized, which you don't want."

For example, say you were to add a metal screw to a metal door. After a while, you would not be able to unscrew it because it would have become part of the door.

Bilén pointed to NASA's Galileo probe, which launched in 1989: lubricant loss and launch vibrations during launch are thought to have stripped the oxide layer from parts of its furled high-gain antenna. When engineers attempted to deploy the antenna in 1991, it never fully opened.

Some metals are more troublesome than others. Gold and platinum don't form an oxide layer at all, even on Earth, which makes them notoriously prone to cold welding. "Gold definitely is a very notorious metal for cold welding," Greer said, adding that gold's softness lets it conform easily to whatever surface it touches and it bonds even more easily.

How to prevent cold welding in space

To prevent components from accidentally fusing in orbit, engineers rely on a few strategies. One is anodizing, a process that locks an artificial oxide layer onto a metal surface. Another method is to coat moving parts with dry lubricants, such as molybdenum disulfide, to physically keep surfaces from touching.

A third strategy is to pair dissimilar metals — for instance, gold next to a "body-centered" metal, like molybdenum — so their atomic structures don't mesh as neatly. "Their packing order is not quite perfectly aligned, and so there'll be a lot more energetic kind of barrier to overcome," Greer said.

Before launch, hardware also gets shaken on vibration tables and cycled through extreme hot-and-cold swings inside vacuum chambers, simulating the stresses of liftoff and orbit to catch problems on the ground.

Even with all of those precautions, cold welding can still happen. Bilén recalled bolts in his own lab's vacuum chamber fusing shut after a move across campus. They eventually had to be drilled out. "It happens even on Earth," he said.

See how much you know about human exploration into space with our spaceflight quiz!

For years, scientists have suspected that, at the molecular level, water is two different liquids ‪—‬ a denser one and a less-dense one ‪—‬ that are constantly switching places. Catching real molecular evidence of this microscopic transformation has been hard. But now, with help from artificial intelligence, researchers say they've finally found it.

"It's hard to imagine — here is just one water, right?" said Xiao Cheng Zeng, a physical chemist at the City University of Hong Kong and co-author of the new study, told Live Science while holding a water bottle in the air. That puzzle sent him digging through scientific literature, where he found the possible explanation: the two-state hypothesis. "That got my attention. We have literature to talk about it but no evidence."

The findings, published June 4 in the journal Nature Physics, could not only prove this long-sought molecular change is real, but also help to explain dozens of water's weird behaviors.

Most liquids become denser as they cool, but water behaves differently; it becomes denser until about 4 degrees Celsius, then starts to expand, which is why ice floats. Water also resists temperature changes better than similar liquids and has a viscosity that decreases under certain pressures. Scientists have documented various anomalies related to water and suspect they may be interconnected.

The two-state model is an attempt to be that unifying explanation.

A 30-year hunch

Zeng has been studying water since his postdoc days in the late 1990s, when he worked on liquid freezing. The two-state hypothesis itself came onto his radar later — around 2006, when he first encountered it at scientific conferences. But for years, he set it aside as too difficult to tackle directly. That changed roughly around 2016, as researchers began reporting experimental evidence that supercooled water could split into distinct high-density and low-density forms.

Around two and a half years ago, Zeng handed the problem to Liwen Li, a postdoctoral researcher in his lab. Rather than repeating the conventional approaches other groups had already struggled with, Li suggested the use of "unsupervised deep learning" — AI trained to spot patterns in data without being told what to look for.

"So AI [is] forced to learn — to use [its] knowledge to create, to explore," Zeng told Live Science.

The team ran massive molecular dynamics simulations, using the GROMACS simulation package. They tracked how hundreds of thousands of water molecules moved and interacted and generated tens of millions of data points.

"Traditionally, you may need a lot of students to figure that out. ... With computers and AI, it took [Li] maybe a year and a half," Zeng said. Without AI, he estimated, the same analysis might have taken closer to a decade.

The AI came back with "reaction coordinates" — a small number of variables, distilled out of all that molecular motion, that describe exactly how a water molecule's local arrangement shifts from the denser structure to the looser one and back. They plotted the system's behavior along those coordinates to see the shape of the conversion. That included the number and location of energy barriers, or saddle points, that molecules have to cross to make the switch.

Two paths up the mountain

The team found that the path the two structures take to convert into each other changes depending on certain conditions. Most of the time, the switch happens along what the researchers call a "semi-loop" pathway, with a single energy barrier to cross.

But near the boundary between high-density and low-density water — the same kind of threshold where ice and liquid water coexist at 32 degrees Fahrenheit (zero degrees Celsius) — the molecules can take a more roundabout "full-loop" path, with three separate barriers instead of one.

Zeng compared it to hiking a mountain that's been sliced in half, with a gentle slope on one side and a sheer cliff on the other. Most hikers stick to the slope; that's the semi-loop. But near the boundary where the two halves meet, it's as if the mountain were becoming whole again, letting hikers circle the entire peak. That's the full loop.

Zeng and his team are now building a more rigorous machine-learning model to confirm the result. They hope to eventually connect it to properties like density, viscosity and temperature.

Confirming the structure in real water won't be simple. Zeng said it will likely require new and sensitive experimental techniques — the kind developed by labs like Pacific Northwest National Laboratory, which previously found indirect spectroscopic evidence for water's two-state behavior.

"Once we have this ... confirmed by experiment," he said, "this model can be used to [understand] how water interacts with nature."

Since most biological and pharmaceutical processes happen in water, a better understanding of water's molecular structure could shed light on how dissolved salts, proteins, and drug molecules interact in solution. "These interactions are vital for injectable drugs and cell function," he noted, but applying this knowledge to practical uses is still a long way off.

Scientists have cooked up a new kind of building material from an ingredient more often found in bread, beer and pizza dough: baker's yeast.

The squishy, yeast-based paste can be squeezed through a 3D printer, dried at room temperature and turned into lightweight architectural pieces, such as wall panels, room dividers and screens that soften harsh sunlight.

Unlike concrete, plaster, or many plastics used in interior decorating, which all use plastic and are difficult to recycle, the new material is designed to minimize waste as it uses renewable ingredients and could eventually draw on industrial leftovers from brewing, agriculture or other yeast-rich processes.

"The interest stemmed from a broader goal of combining circularity-oriented architectural design, sustainable biomaterials, and digital fabrication to develop a fully bio-based architectural material from abundant, renewable resources," Malgorzata Zboinska, a professor of architecture at Chalmers University of Technology in Sweden and an author of the study, told Live Science in an email.

The study was published on March 5 in the journal Frontiers of Architectural Research.

Building with yeast

To make the biomaterial, the researchers first heated up the yeast to deactivate it, so it would not be alive in the finished product. They then mixed it with wood-derived cellulose fibers, algae-derived gel, called alginate, plant-based sugars and water. The mixture resulted in a smooth hydrogel, a soft jelly-like material that can hold a specific shape and be molded with a 3D printer.

"We use pressure-based 3D printing at room temperature, which is important due to its sustainable aspects — it does not require energy-intensive heating or additional support structures," Zboinska said.

After printing, the pieces were left to dry in room temperature conditions. As the water left, the gel stiffened into a stable, lightweight solid. The strongest versions reached an average tensile strength of 2.7 megapascals (391.6 psi) — around the strength of a fruit roll-up (or fruit leather) — and stretched up to 25.2% before breaking. While the material isn't that strong, it is effective at holding its shape, which is key for making products like screens and wallpaper.

"Structurally, we found that yeast contributes differently depending on how it is processed," Zboinska said. "This allows us to tune the material’s properties through relatively simple formulation changes."

The researchers found that when the yeast cells remain intact, they act mostly like a filler, giving the material volume. But when the yeast is deactivated, they release internal components that help bind the mixture together.

By changing the recipe and printing pattern, the team could alter the material's color, texture, porosity and translucency. In the study, printed prototypes measured 7.87 by 19.69 inches or (20 by 50 centimeters), and let through between 5.6% and 31.6% of light, depending on their design.

A greener future for interior designers

The construction sector uses huge amounts of raw material and energy, and researchers are hunting for lower-impact alternatives. Zboinska and her team hope yeast-based materials could replace some fossil fuel-derived interior products, like synthetic tiles, drapes or plastic panels, rather than load-bearing materials like steel and concrete.

"Biomaterials are … commonly viewed as safer for the environment upon disposal," Timothy Long, center director and professor for the Biodesign Center for Sustainable Macromolecular Materials and Manufacturing at Arizona State University, who wasn't involved in this study, told Live Science via email. Long cautioned that even if biomaterials like this yeast-based product are designed to minimize waste, they only work if protocols are in place to help with proper disposal.

"So even though they are biobased materials, we need to focus as a community to collect, recycle, and reuse these materials," he said.

Yet, Long believes that even if there aren't proper recycling practices for these special materials, they can still have a positive impact on the environment.

"There is also evidence that if biomaterials remain in a biological environment then their decomposition products are more likely to be safer to humans and safer for the Earth" than non-biodegradable materials, he said.

Still, for the yeast-based material, big questions remain. The team has not tested how long the material lasts, how it handles moisture over time or how it behaves thermally or acoustically. They have also not explored whether the deactivated yeast could trigger reactions in people with yeast allergies.

Before the material could move into real buildings, the researchers would also need to improve printing precision, scale-up methods and further fine tune how the material bends and shrinks as it dries, Zboinska said.

But for now, the work suggests that the future of interior design could begin with a vat of humble yeast.

"The research points toward new ways of thinking about circular design and sustainable manufacturing in architecture, where fabrication processes, material behavior, and environmental considerations are closely integrated from the outset," Zboinska said.

Scientists have invented a self-cleaning coating that prevents stains from food, oil or dirt from sticking to a fabric's surface, making the spots easy to remove without detergent or large amounts of water. The coating, which creates a protective ultrathin layer of water its developers call "molecular water armor," could reduce the water and electricity demand of household laundry by over 80%, new research suggests.

Washing clothes is a basic necessity of everyday life, but laundry has a huge environmental footprint because it produces vast amounts of wastewater. A single cycle in a household washing machine uses 10.5 to 16 gallons (40 to 60 liters) of clean water, while the detergents required to remove stains promote the release of microplastics from synthetic fibers and leave chemical residues in the discharged water.

Better detergents and more efficient washing machines are slowly reducing the ecological impact, but with an estimated 2.6 billion gallons (10 billion liters) of wastewater generated annually from laundry in China alone, there's a need for additional measures.

In the new work, reported in the journal Communications Chemistry March 19, researchers tackled this problem from a completely different perspective, focusing on the design of the textiles themselves.

"Instead of relying on detergent to remove strongly attached dirt, we modify the textile surface so that stains do not adhere strongly in the first place," material scientists Chongling Cheng of Southeast University and Dayang Wang of Jilin University in China, both co-authors of the research, told Live Science in a joint email. "As a result, many common stains can be removed using water alone and a much shorter rinsing process."

"Molecular water armor"

The team, based at Southeast and Jilin universities in China, spray-coated alternating layers of positively and negatively charged polymers onto cotton, silk and polyester clothing. The resulting multilayer film formed a surface rich in sulfonate groups — sulfur-containing chemical units — which attracted and organized water molecules into an ultrathin layer on the fabric's surface.

"We sometimes describe this layer as a molecular 'water armour,'" Cheng and Wang said. "[It] acts as a barrier between the fabric and contaminants. Oils, food stains, sweat residues, and microorganisms find it difficult to make direct, strong contact with the coated fiber surface. Therefore, stains are much less strongly attached and can be removed by water flow, without the need for detergent." This contrasts with waterproof materials which simply allow water droplets to roll straight off, rather than producing a cleaning effect.

The researchers demonstrated this superior cleaning performance by pitting the new coating against conventional detergent-based washing to launder cloth soiled with ketchup, chili oil and soy sauce. It proved particularly effective for these oily surface stains, and the coated fabrics cleaned with just a single rinse — matching or outperforming the detergent wash for stain removal.

An analysis of the small amount of wastewater produced by the rinse wash also revealed that this additional layer drastically suppressed the release of microplastics during washing, trapping any stray particles within the coating itself.

Overall, the polymer layer substantially improved the environmental footprint of the laundry process; the team calculated that it reduced the water and electricity demand by more than 82% compared with a normal cycle. From a cost perspective, these stain-resistant fabrics are more expensive to manufacture than standard textiles, but the team noted that this upfront expenditure could be recouped in as few as 15 laundry cycles, depending on the cost of the detergent.

The coating also addresses laundry hygiene. The coated fabrics showed both antibacterial and antifungal effects, which are most likely related to the water armor barrier, Cheng and Wang said. The repellent surface prevents microbes, sweat and skin cells from adhering to the fabric, allowing them to be easily removed by rinsing. As a result, a quick soak completely eliminates odors and prevents the growth of mold or mildew, which can occur during prolonged storage.

But convincing consumers of these clean credentials could prove challenging. "Many people associate detergent foam and fragrance with cleanliness, even though foam and fragrance are not necessarily direct measures of hygiene," Cheng and Wang said. "Therefore, we expect that consumer trust will require clear evidence, transparent testing, and practical demonstration."

The coating is still in the proof-of-concept stage. Preliminary investigations by the team have suggested that it is safe to use against skin and remains effective over more than 100 laundry cycles. Given the coating's nanoscale thickness, the team doesn't expect it to change the texture or breathability of treated clothing. However, a crucial next step will be to have consumers test the materials to validate their comfort and practicality in real-world settings.

"Commercialization will also require independent safety assessment, durability standards, [and] environmental lifecycle analysis," Cheng and Wang said. "Our current study demonstrates the central scientific principle. The next challenge is to translate this principle into a robust, affordable, and trusted product for everyday use."

On a dark July morning in 1945, U.S. scientists and military personnel detonated the world's first nuclear bomb in a remote area of New Mexico. The blast unleashed the energy equivalent of 25,000 tons of TNT, completely vaporizing the bomb's drop tower and reducing the desert sand within a 1,000-foot (300 meters) radius to glass.

Scientists later dubbed this pale-green-and-red, faintly radioactive glass "trinitite" after the test site, Trinity. Now, more than 80 years later, researchers have discovered that some red trinitite contains unique crystals found nowhere else in nature. They detailed the finding in a study published May 11 in the journal PNAS.

The investigation was sparked by a different mineral: an unusual quasicrystal previously identified in samples of red trinitite. Unlike most quasicrystals, which are composed primarily of aluminum, this quasicrystal is rich in silicon. Its existence suggested there might be other bizarre crystals hidden in the Trinity glass.

"We wanted to further explore these extreme-formation products," Luca Bindi, a mineralogist at the University of Florence in Italy and first author of the new study, told Live Science in an email.

History in a crystal

Bindi and his team used an electron microprobe and X-ray diffraction to examine a rare "oxblood" variant of red trinitite. The striking crimson color of this sample came from the disintegrated test tower and the metal equipment surrounding it. Metallic droplets from these structures were trapped inside the molten silicon glass as it fused in the blast, changing its hue from sage to scarlet.

In this sample, the researchers found a never-before-seen clathrate crystal. Clathrates are a type of crystalline structure in which one element forms a "cage," trapping other atoms inside. In this case, atoms of silicon enclosed copper and calcium inside linked 12- and 14-sided crystal lattices. This type of arrangement is rare in nature, especially for inorganic compounds, the team said.

This marks the first time clathrate crystals have been found as a byproduct of a nuclear blast. During the Trinity explosion, temperatures exceeded 2,700 degrees Fahrenheit (1,500 degrees Celsius), and pressures briefly climbed to 8 gigapascals — comparable to the pressure deep beneath Earth's crust. Such intense conditions forced atoms into configurations they normally wouldn't be able to take.

The team also investigated the possibility that the new clathrate may have been a precursor to the previously described trinitite quasicrystals. A mathematical analysis showed that this was unlikely. But exploring this relationship helps fill out our knowledge of the upper limits of mineral formation, well beyond anything that can be replicated inside a lab.

"Extreme events like nuclear blasts, lightning, or impacts can generate new mineral phases and structures that expand our understanding of how matter organizes under extreme conditions," Bindi said.

Scientists have developed a method to turn plastic waste into clean hydrogen using solar power and acid from old car batteries.

The one-pot process transforms hard-to-recycle plastics into valuable industrial chemicals and clean fuel, potentially creating a circular upcycling system that tackles multiple problematic waste streams at once, the researchers say.

The world produced more than 440 million U.S. tons (400 million metric tons) of plastic waste in 2025, but less than 10% of this was actually recycled. The challenge lies in the sheer number of different plastics we use. Some, such as polypropylene and polyethylene, can be easily melted down and remolded, while others require specific chemical processes to break the polymer structure down into its individual chemical building blocks (known as monomers).

Condensation polymers like polyethylene terephthalate (PET, often used for packaging food and drinks), polyurethane (PU, which can be used in foam cushioning, bedding and insulation), and nylon fall into this latter category. A chemical reaction between two different monomer units releases water to form bonds between these fragments, creating a long alternating polymer chain. These bonds can later be broken by adding water back to the molecule, releasing the monomer building blocks and breaking down the plastic.

In the new study, researchers took this one step further — not just recovering the monomers but also upcycling the plastic waste into other valuable chemical products.

The team set its sights on hydrogen, a green fuel source and an important industrial feedstock, and developed a process to combine plastic depolymerization and hydrogen generation in a single reactor. While both steps have been studied individually before, no one has ever achieved them together. The researchers reported their findings in the journal Joule April 6.

The scientists began with the depolymerization step. Focusing on PET, they ground samples of plastic bottles into a fine powder and dissolved them in concentrated sulfuric acid. "We heat that up to 140°C [degree Celsius, or 284 degrees Fahrenheit] and that hydrolyses the plastic back into its monomers," study first author Kay Kwarteng, a researcher at the University of Cambridge, told Live Science. "For PET, that is ethylene glycol and terephthalic acid," which are both valuable industrial chemicals, he added.

However, rather than using fresh sulfuric acid from a bottle, the team saw an opportunity to harness another problematic waste stream. "Sulfuric acid is a component of car batteries, but when they are recycled, they only recover the lead component," Kwarteng said. "We could extract the battery acid and use that instead. It makes a strong argument for sustainability."

The terephthalic acid conveniently precipitates out of the reaction as it forms, leaving an acidic mixture rich in ethylene glycol.

However, the second step, which produces hydrogen from the ethylene glycol monomer, usually needs alkaline conditions to work. The sunlight-powered reaction breaks the ethylene glycol down into even smaller chemical products, but the researchers first had to design a new catalyst that would remain stable in the battery acid.

They settled on a molybdenum metal system and added it directly to the mixture. "Once we expose the catalyst to light, it oxidizes the ethylene glycol which generates electrons," Kwarteng said. "These electrons can convert protons," — present in the acid mixture — "to hydrogen, and they oxidize the ethylene glycol to acetic acid."

The hydrogen and acetic acid formed in this process are less valuable than the ethylene glycol monomer, but crucially the approach provides a sustainable entry point for other related chemistry, said Erwin Reisner, professor of energy and sustainability at the University of Cambridge. "Instead of making hydrogen, we can hydrogenate organics," he told Live Science. "It's exactly the same system, but instead of evolving hydrogen, we just add unsaturated organics and hydrogenate them directly."

Hydrogenation is an important industrial reaction that inserts hydrogen across a double bond, typically using hydrogen generated from fossil fuels. But in a follow-up study published in the journal Angewandte Chemie International Edition on Monday (May 4), the researchers demonstrated how their new process could be used to hydrogenate nitrogen-containing substrates into important pharmaceutical building blocks. "When we use plastics for this hydrogenation, we reduce the carbon footprint by half," Kwarteng said.

The team are now looking at tailoring the reaction design for the needs of industry and plan to test the process in a flow reactor — a system which continuously converts reactants to products, rather than producing hydrogen in batches.

The use of so many recycled reagents is impressive, Amit Kumar, a catalysis researcher at the University of St Andrews' School of Chemistry, told Live Science. But he noted that the photochemical step could prove challenging for industry. "I think it's super interesting that you can just use this plastic as a hydrogen source and science-wise it's very exciting that you can use visible light," he said. "The next step towards commercialization will be scaling up and demonstrating the process in flow."

"Forever chemicals" — technically known as perfluoroalkyl and polyfluoroalkyl substances (PFAS) — are remarkably useful things. They're also among the most dangerous pollutants on the planet, says investigative journalist Mariah Blake.

The many carbon-fluorine bonds in these chemicals, consisting of an alkyl chain connected to multiple fluorine atoms, are considered the strongest in organic chemistry.

This makes them remarkably resistant to heat, water, and chemical corrosion, giving them applications in everything from nonstick cookware to firefighting foam, medical devices and waterproof clothing. In fact, their first major application after their invention was in separating weapons-grade uranium isotopes from corrosive gas for the Manhattan Project, assisting in the creation of the first atomic bomb.

But as Blake recounts in her book "They Poisoned the World" (Penguin Random House, 2025), PFAS bonds are "proteinophilic", causing them to strongly bind with proteins in the organs and blood, persisting there for years. Outside of our bodies, the chemicals have biological half lives that can span centuries. Eight decades since their invention, they are now in the bodies of nearly every human being on the planet — accumulating in our bloodstreams, livers, kidneys and lungs.

Blake's book chronicles a decades-long cover-up that hid the chemicals' links to cancer and birth defects, and the fight to expose the contamination by the people of Hoosick Falls — a village in New York State whose water was polluted by PFAS runoff. Her work has been shortlisted for multiple prizes, including the PEN/E.O. Wilson Literary Science Writing Award, an annual award for excellence in nonfiction in the physical or biological sciences.

Live Science spoke with Blake about PFAS, how they came to pollute the planet, and why one of history's greatest corporate scandals is not quite over yet.

Ben Turner: Let's start off simple. What are PFAS?

Mariah Blake: PFAS, also known as forever chemicals, are a large family of substances with some pretty incredible properties that make them very useful.

They're extremely resistant to heat, stains, water and grease. They stand up to corrosive chemicals that burn through almost every other substance. They helped usher in air and space travel and high-speed computing. They have given rise to lifesaving medical devices — things like patches for deteriorating veins and arteries. They've transformed thousands of everyday items — everything from dental floss and clothing to kitty litter and makeup. They are in all of our homes and in all of our blood.

They also happen to be, in my opinion, the most insidious pollutants in all of human history.

They persist in the environment for hundreds, or even thousands, of years. Those that have been studied are highly toxic, even in the most minuscule of doses, and they are literally polluting the entire planet, including human blood and ecosystems in the remotest parts of the world — so places like the Tibetan Plateau or Mount Everest or the deepest parts of the ocean. They are literally everywhere.

BT: PFAS were first developed in 1938 with the invention of Teflon, but they weren't immediately used for commercial purposes. In fact, you write that their first major use was in the Manhattan Project, the secret project led by J. Robert Oppenheimer to develop the atomic bomb.

MB: There had been a couple of PFAS that were developed prior to the war as a result of laboratory accidents, but they never would have been produced on a commercial scale if it weren't for this U.S. government program. The U.S. government had physicists working in labs all across the country to develop nuclear fuels and the bomb itself, and they had chemists working to develop PFAS. They developed various methods to produce them, and they put them into mass production as early as 1943.

It was clear from the beginning that these were dangerous chemicals. The plants where they were manufactured were prone to fires and explosions; workers were regularly hospitalized with breathing problems and chemical burns or worse.

But it wasn't just workers who were affected. Beginning in 1943, farmers downstream of these plants began to complain that their peach crops were burning up, that their cows were so crippled they couldn't stand, and they had to graze by crawling on their bellies.

They began to complain to DuPont [which manufactured Teflon or Polytetrafluoroethylene using Perfluorooctanoic acid (PFOA), a synthetic chemical in the PFAS family] because nobody knew that the Manhattan Project was happening.

These complaints alarmed Manhattan Project officials, and they launched an elaborate research program to study the health and environmental effects of the chemicals. They had determined as early as 1947 that PFAS were highly toxic and that they were accumulating in the blood of people around the [chemical] plants.

The FDA [Food and Drug Administration] also began studying the fluoride content. They didn't have the technology then to detect these specific chemicals, but they started testing the [food] produced around the plants for fluorides known to be toxic at high levels. They determined that there were extraordinarily high levels in the produce, and they were going to ban the produce from this region.

But Manhattan Project officials intervened to stop that from happening. And I think that's really key, because if that had happened, concern about that would have raised a red flag. People would have begun looking into these chemicals much, much earlier. So that decision set public knowledge back by half a century or more.

BT: Instead, in the postwar commercial boom that followed, chemical companies began to mass-produce PFAS — for cookware, fabrics, food packaging, in cars, planes and industrial processes. Then the chemicals started leaking out through landfills and industrial discharge to waterways, and later into our bodies. Which companies were responsible? And how early on were they aware that their products were toxic?

MB: After the war, a Minnesota-based company called 3M [originally the Minnesota Mining and Manufacturing Company] acquired patents for technology to produce PFAS. They actually hired some Manhattan Project chemists to parlay them [PFAS] into substances that could then be marketed to corporations for manufacturing and to the general public. So as early as the 1950s, you had products like Scotchgard and Teflon appearing on the market that incorporated these chemicals.

Now, the industry was also aware very early on that these chemicals were harmful. As early as the 1960s, the two main manufacturers, DuPont and 3M, knew that they were toxic. By the 1970s, they had discovered that these chemicals were accumulating in the blood of people all over the United States, even in places where there was no known source ‪—‬ so places where they weren't being used in manufacturing. They eventually discovered that this was true all over the world.

They looked at thousands of blood samples collected from around the globe, and some from past medical studies. The only samples they could find anywhere that didn't contain these chemicals were collected from Korean War veterans before 1952 — before these chemicals went into wide-scale production. What that tells us is that these chemicals were probably already ubiquitous in the environment by the 1960s.

The blood data set off alarms inside DuPont and 3M, and they began intensively studying the health and environmental effects of these chemicals. They quickly discovered that they [PFAS] did not break down in the environment at all and that they had a devastating effect on lab animals.

In one case, they tested the Teflon chemical PFOA [perfluorooctanoic acid] on monkeys, which were chosen because they're more biologically similar to humans than lab rats, and all of the monkeys died.

These two companies [3M and DuPont] also began monitoring the effects of the chemicals on their workers, and they linked them to a number of diseases, like kidney cancer, prostate cancer, leukemia, organ damage, drops in testosterone, and immune suppression.

But most alarmingly of all, they found that there was a link to birth defects. In the 1970s, 3M conducted a study that found that rats who were exposed to this Teflon chemical [PFOA] while they were pregnant gave birth to pups with facial deformities.

DuPont decided to see if the same was true of its workers, conducting what it called a "pregnancy outcome questionnaire". The goal, according to internal DuPont documents, was to determine whether this chemical caused "abnormal children."

Two of eight women who gave birth during the course of this study gave birth to children with facial defects very similar to the ones that had been found in rats. I interviewed one of these women [Sue Bailey] and her son [Bucky Bailey] when I first began reporting on this story, and their experience was completely wrenching. This young man went through 40 or so surgeries within the first year of his life to correct these deformities.

BT: So what happened to the evidence?

MB: Rather than alerting the public or regulators, DuPont simply shut down the study and continued exposing workers to these chemicals. In fact, they temporarily moved female workers out of the area where they were exposed to these chemicals and then moved them back.

But I think, perhaps more importantly, they continued exposing the public to these chemicals.

By this point DuPont was aware that these chemicals were polluting drinking water around its plants and all over the country. And there are very simple steps that DuPont could have taken to filter these chemicals so they didn't leave its factories, but it opted not to do so because the additional expense wasn't justified in the eyes of executives.

BT: That seems shortsighted, given that it has already led to big lawsuits. Are you suggesting they thought they could get away with it?

MB: They came to the conclusion that they would already be liable for the 32 years that they'd been producing the chemical, and so any additional liability would be incremental.

It was a very cold calculation. They didn't think that they wouldn't get caught; they just thought that the additional expense of installing these filters would be less than the additional liability they would face from not installing them.

BT: The thing you note that really puts the cherry on all of this is that tests have found there really is no safe level of exposure to these chemicals.

MB: They couldn't find a dose at which PFAS didn't have health impacts.

The U.S. EPA [Environmental Protection Agency] has set safety standards for the two best known and best studied of these chemicals [PFOA and PFOS] in drinking water, and the safety standard is 4 parts per trillion — which is the lowest level you can reliably detect — but the health based goal is zero. So the EPA has essentially said there is no safe level of exposure to these chemicals.

BT: In the end, it was left to the ordinary people in the worst-hit places to fight back. How did regulators — the EPA, the FDA, even mayors of poisoned towns, look the other way for so long?

MB: Most of the evidence was not made public, right? Industry withheld all of its own internal data. And they were allowed to do this partially because of the way we regulate chemicals.

When the current [chemical] regulation [system] in the United States was developed, chemicals that were already on the market were presumed safe and grandfathered in.

Things are a little different in Europe, which has since embraced the precautionary principle [Editor's note: The precautionary principle means regulators can ban or restrict a substance if there's some evidence of danger to human health, even if this link has yet to be proven].

But as a result of this, the vast majority of chemicals [on the U.S. market] have never been tested for safety, and companies don't have to provide their own internal data.

In fact, the only reason PFAS even came onto regulators' radar ‪—‬ and later, the radar of scientists and the public ‪— was because a family of West Virginia farmers sued DuPont after runoff from a DuPont landfill began killing off their cattle. That exposed this whole cover-up and is the only reason that the world knows these chemicals exist.

But it didn't happen until the late 1990s. So scientists, the public and regulators weren't aware that these chemicals existed — despite having been in circulation for 80 years — until about 25 years ago. That's a key piece of the puzzle.

But even after this information came to light, the chemical industry launched this Big Tobacco-style campaign to downplay or suppress evidence that these chemicals were harmful. They deployed all kinds of strategies: They published their own internal studies in the scientific literature with more troubling findings downplayed; they hired outside scientists to defend the safety of these chemicals; they funded or founded ostensibly independent think tanks and scientific organizations to discredit the science on PFAS and other toxic chemicals, and undermine science-based regulation; and they flooded Congress with lobbyists.

As a result, even after this information entered the public record, it took another 15, almost 20, years for it to really register with the public.

BT: You write in your book that, more recently, chemical companies have switched to using shorter chain fluorocarbons. What are they? How widely are they being used? And how dangerous are they compared to legacy chemicals like PFOA?

MB: The truth is that we don't know what chemicals are being used, and how widely they're being used, because manufacturers don't have to disclose which chemicals they are using. So when we have discovered what chemicals are being used in place of the two that have theoretically been phased out, it's more or less by happenstance or because scientists go through some elaborate detective work.

Now we do know that PFOA, the Teflon chemical that I've spoken about, was replaced in DuPont's [whose performance chemicals business was later spun out to become Chemours] formulations, at least, with Gen X. DuPont claimed that Gen X was safer than PFOA, but we now know that that is probably not true. So the evidence that has been collected so far shows that Gen X is probably as toxic as PFOA, and in other ways it is more problematic as it moves more readily through the environment.

They [shorter-chain fluorocarbons] build up more rapidly in crops, leading to higher concentrations in food, and they are more difficult to remove from drinking water. All over the country, communities have spent hundreds of millions or even billions of dollars installing filtration systems to remove the better studied PFAS.

But [these systems] aren't particularly useful for Gen X. And scientists are now discovering that there are other forms of PFAS that are even more abundant in the environment that can't be removed with the existing technologies at all, at least not technology that's feasible to deploy on a utility scale.

One of the EPA scientists I spoke to in the course of my research likened it to cutting off the head of the Hydra and having it sprout more to replace it. He was one of the scientists that spent years reverse-engineering the identity of Gen X by testing water downstream of a plant.

BT: What can we do to finally stop PFAS? Especially if harmful new chemicals are developed quicker than they can be detected and identified, let alone regulated?

MB: You can't as long as you're regulating chemicals one by one. But if you start regulating them as a class, I think it's possible for regulatory agencies to be much more effective. In the case of PFAS, you have a class of probably around 15,000 chemicals.

And I think in response to the unique threat they pose, you actually do have some government agencies beginning to regulate them as a class. A lot of the U.S. states that have bans have banned the entire class of chemicals, and they include exceptions for uses that are essential to the health, functioning and safety of society, and for which there are no substitutes available. But otherwise they are banned. End of story.

The EU ban is supposed to be a class-wide ban. There are lobbyists in Brussels right now fighting to insert loopholes into it, but I think Europe has a historic opportunity right now, because a strong PFAS ban in Europe would move us a lot closer to turning off the tap on these chemicals. Regulating them as a class is the only solution.

Editor's note: This interview has been condensed, and edited for clarity. 3M did not respond to Live Science's request for comment by the time of publication.

DuPont de Nemours told Live Science in an email that it emerged as a new and independent company in 2019, and that it cannot comment on products, events and other actions that occurred while E.I. DuPont de Nemours ran its performance chemicals business, or after that business was spun off into the independently-owned and operated Chemours Company in 2015.

Breadcrumbs from food waste could replace fossil fuels as a source of hydrogen in one of the most common chemical reactions used in chemical manufacturing, new research suggests.

The new process, reported Feb. 23 in the journal Nature Chemistry, combines natural fermentation processes in bacteria with metal catalysis to generate an array of valuable chemical products from simple food waste. Calculations showed that this hybrid procedure was carbon negative overall, and the authors think it could be the first step in reimagining chemical manufacturing as a more sustainable industry.

Hydrogenation is a chemical process that inserts a molecule of hydrogen across a double bond and is a staple reaction in food production, plastics manufacturing and the synthesis of drug compounds.

However, the bulk of the hydrogen gas used in this reaction is derived from fossil fuels through a dirty and energy-intensive process called steam reforming, which produces 15 to 20 kilograms of carbon dioxide for every kilogram of hydrogen generated . Consequently, hydrogenation is a huge sustainability challenge for the chemical industry, and scientists are urgently searching for greener alternatives.

Turning to nature, Stephen Wallace, a professor of chemical biotechnology at the University of Edinburgh, decided to investigate whether it was possible to harness the power of biology to tackle this chemistry problem. Many bacteria naturally produce hydrogen when they are forced to respire anaerobically (without oxygen), and they release a constant stream of this gas into their surroundings. If this could be linked to a compatible chemical system, it would be theoretically possible to use bio-hydrogen in a hydrogenation reaction, thereby eliminating the need for fossil fuels in this process, Wallace reasoned.

"The main challenge was finding a catalyst that can operate in a living system ‪—‬ in water, at mild temperatures, and without harming the cells," he told Live Science in an email. "We had to balance both sides: a catalyst that stays active in a complex biological environment, and microbes that continue functioning in the presence of the catalyst."

Culture shift

The team cultured E. coli bacteria in a glucose-containing medium, adding a commercial palladium catalyst and a test substrate before sparging the mixture to remove oxygen. The oxygen-free reaction was incubated at 98.6 degrees Fahrenheit (37 degrees Celsius) for a day, and subsequent analysis revealed that the top-performing strain had produced the expected hydrogenation product in 94% yield.

"The metal catalyst comes in and is essentially bound to the cell membrane," Simone Morra, a biotechnologist at the University of Nottingham who wasn't involved in the work, told Live Science. "The cell itself will produce the hydrogen, and then as soon as the hydrogen starts to diffuse out of the cell, it will hit this metal catalyst, which will do the second part of the reaction and produce a hydrogenation product."

With a biocompatible system established, Wallace next sought to replace the expensive glucose feedstock with a cheaper and more sustainable alternative. Focusing on bread waste, the team used microbial enzymes to break the complex carbohydrate molecules within breadcrumbs into simple glucose units. This waste-derived fuel was then fed directly to the E. coli cultures, effectively converting breadcrumbs into hydrogen.

But the researchers had one final trick up their sleeves: Instead of feeding a precursor molecule to the bacterial culture, they genetically engineered certain strains to produce the required substrates within the cells themselves. "It’s brilliant and very inspiring," Morra said. "They show that they can capitalize on the synthetic abilities of E. coli. Essentially they can make use of the carbon pathways of the cell to make any substrate they want."

The use of bio-generated hydrogen resulted in a three-fold decrease in greenhouse gas emissions compared to using fossil fuels. The breadcrumb-powered hydrogenation process, in particular, reduced the global warming potential by more than 135%, corresponding to a carbon-negative footprint.

The team is now working to increase the number of possible substrates and developing the process to accept more types of biowaste. Ultimately, they hope the method could be incorporated into industrial chemical synthesis.

"Right now, the system works best with simpler alkenes," or molecules containing a carbon-carbon double bond, Wallace said. "It's not yet as efficient as industrial processes, but it demonstrates a fundamentally new way of doing hydrogenation. To make it viable, we need to improve efficiency, scale the biology, and develop catalysts that remain stable and cost-effective at industrial scale."

A college student has created a prototype polish to turn a fingernail into a touch-screen stylus, after noticing that people with long nails and calloused fingertips struggled to work their smartphones.

Manasi Desai, a student at Centenary College of Louisiana with an interest in cosmetic chemistry, launched the project with her research supervisor, Joshua Lawrence, an associate professor of chemistry at Centenary. Their goal was to create a clear, nontoxic polish that would allow a nail to access a touch screen the way a human fingertip does.

"Our final, clear polish could be put over any manicure or even bare nails, which could help people with calluses on their fingertips too," Desai said in a statement. "So it has both a cosmetic and lifestyle benefit."

Desai and Lawrence presented their research Monday (March 23) at the annual meeting of the American Chemical Society.

Who needs touchscreen polish?

While touch screens are nearly ubiquitous today, some people cannot use them as easily as others. For example, guitar players and carpenters with callouses on their fingers may be unable to get the screen to register their touch because their fingertip skin prevents electrical flow ‪—‬ a problem Consumer Reports called "zombie finger" in 2015. But touch screens are also difficult to use when a person's hand is gloved, is very dry or has long nails.

After the researchers asked a phlebotomist with long nails who was struggling to operate a smartphone whether touch-screen-compatible nail polish would be useful — and got a resounding "Yes" answer — they set out to develop a new product.

"Chemists are here to solve problems and to try to make your world better," Lawrence said in the statement.

The touch screens in modern tablets and smartphones work through a property called capacitance. A small electric field is created on the screen, and when a conductive material, like a finger, touches the surface, the electric field is interrupted. The screen registers the disruption — a change in capacitance — as a touch at a particular location. But tapping the screen with a nonconductive material, like a fingernail, does not register as a touch.

Previous researchers' attempts at creating a capacitive nail polish focused on including carbon nanotubes or metallic particles to make the fingernail electrically conductive. However, these particles are dangerous if inhaled and limit the color range of polishes.

For her project, Desai methodically tested combinations of 13 commercially available clear-coat nail polishes and more than 50 additives to find one that met three criteria: It was clear, it was nontoxic and it created a conductive top coat.

In her experiments, Desai found that the polishes that performed best included the amino acid taurine and the organic molecule ethanolamine, an amino alcohol. When combined, the taurine and ethanolamine additives created a formula that registered as a touch on a smartphone.

The new polish formula was designed to work through acid-base chemistry rather than the inclusion of metal additives. With acid-base chemistry, acids donate protons and bases accept protons.

"We think that the materials we are producing are working via protons hopping from acidic to basic groups," Lawrence told Live Science in an email, because the mixture of taurine and ethanolamine — an ammonium acid and an amine base — works well. "We think we have proton exchange between acidic and basic groups at the surface of the polish, fulfilling the same role as ion mobility in skin," Lawrence said.

It will be a while before their polish hits store shelves, though. The nail polish does not work long enough yet. "All our formulations lose efficacy too quickly," Lawrence said. "They stop working after hours or days, and we want them to work for days or weeks, minimum."

Desai and Lawrence are working on tweaking their formula to find the best-performing combination of ingredients and to make the current formulation nontoxic. Currently, the least-toxic formulation they devised results in a gritty, speckled finish ‪—‬ "not high fashion to be sure," Lawrence said.

The researchers have already submitted a provisional patent for their invention. "Right now, we have a good proof of concept material, but need to do a lot more work!" Lawrence said.

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

Researchers have created an unusually twisted molecule with a never-before-seen electronic structure.

The new molecular architecture, dubbed half-Mӧbius topology, "is another knob that we can turn in order to make and manipulate matter," and expands our fundamental understanding of physics and chemistry, co-lead author Igor Rončević, a lecturer in computational and theoretical chemistry at the University of Manchester in the U.K., told Live Science.

A Mӧbius strip, which is created by twisting a ribbon 180 degrees and then joining the ends, is a mathematically interesting shape that results in a single continuous surface. This weird inverted geometry also has interesting implications for chemists, particularly when they're considering the electronic and spatial properties of molecular structures.

Electrons in revolt

Usually, electrons are localized around a specific atom or bond, but a subset of cyclic compounds, known as conjugated rings, allow the electrons to travel freely throughout the entire loop, above and below the atoms. This delocalization makes conjugated rings more stable than expected, and also influences other properties, including color, optics and reactivity.

However, in a Mӧbius molecule, the electronic orbitals holding the electrons are twisted 180 degrees relative to each other at the junction where the ends meet. The electrons can still move across the whole molecule, but at this junction, some of their properties effectively cancel out, resulting in completely contrasting characteristics and behavior for the overall molecule.

"Chemistry thought that these are the only two options," Rončević said. "But our discovery shows that there's another option, a third option, where we can also rotate by just 90 degrees."

To achieve this, the team, co-led by Leo Gross, principal research scientist at IBM Zurich, created two conjugated systems within a single ring of 13 carbon atoms. The ring contained two chlorine atoms bonded at positions 1 and 7 which isolated these conjugated systems and unevenly separated the electrons on each side. One side of the ring held 13 electrons, while the other held only 11.

"The problem is, electrons like to pair up," Rončević said. "So what they will do in order to pair up is, they will twist the molecule."

The ring, therefore, spontaneously twists itself by 90 degrees ‪—‬ pushing one chlorine atom up and the other down ‪—‬ to align these two separated conjugated systems. This then enables mixing between the two systems, allowing them to share their electrons across the whole molecule.

"At this point, we don't have two separate systems any longer; we have one 24-electron system," Rončević said. The resulting molecule therefore exhibits its own characteristic electronic and magnetic properties, distinct from both standard and Mӧbius structures.

One last twist

The half-Mӧbius molecule's restricted twist angle also results in two possible versions of itself, known as enantiomers.

Because the ring can twist either left or right, the resulting molecules are mirror images of each other ‪—‬ much like left and right hands. This property, technically called chirality, is hugely important throughout chemistry, affecting everything from the synthesis of drug molecules to the production of OLEDs. Intriguingly, by applying a small external voltage the team could freely interconvert a single molecule between the two enantiomers — something that is immensely difficult to achieve using conventional chemistry.

The team supported these experimental findings with detailed computations; the mind-bending complexity of the half-Mӧbius electronic structure necessitated state-of-the-art quantum computers. They published their findings March 5 in the journal Science.

Looking forward, the team intends to focus on exploring the fundamental theory and potential of these molecular architectures.

"We really made a molecule that has a completely new electronic structure, and we want to see what else is possible," Gross said. "We could expand this and explore, for example, several half-Mӧbius twists or even braided ones."

Researchers in China have made what they claim to be the first samples of pure hexagonal diamond, a theorized rare variant of superstrong diamond found in meteorites from shattered dwarf planets.

Natural diamond, also called cubic diamond, has been considered the hardest natural material on Earth for so long that the Mohs hardness scale, which rates minerals' resistance to scratching, uses diamond as the scale's upper limit. It's called cubic diamond for its neat arrangements of carbon atoms in a cubic structure. In contrast, hexagonal diamond organizes carbon atoms in a lattice made of hexagons, like a honeycomb.

An elusive mineral

In 1962, researchers at the Pittsburg Coal Research Center theorized that layers of carbon atoms making up diamond could be organized in a hexagonal lattice instead of a cubic one, thanks to how carbon forms bonds with other carbon atoms. In 1967, researchers discovered hexagonal diamond — or lonsdaleite — in the lab, suspecting it could be harder than cubic diamond.

They started looking for it in a special type of diamond-rich meteorite called ureilite, which forms from the mantle of smashed dwarf planets. The first detections of hexagonal diamond in the wild were documented in a 1967 paper; three Canyon Diablo meteorites (fragments of an asteroid that created a large crater in Arizona) with about 30% hexagonal and 70% cubic diamond phases, and Goalpara meteorites (found in Assam, India) that had a small amount of hexagonal diamond.

Not everyone agrees that the Canyon Diablo lonsdaleite exists. Some scientists thought the evidence could be explained by flawed cubic diamond that was stacked chaotically, and they weren't convinced that lonsdaleite had been detected in previous studies. However, multiple recent studies have identified lonsdaleite in meteorites and in lab samples, including a 2025 study that made small amounts of it in the lab.

The biggest challenge in identifying lonsdaleite is the lack of pure samples; in many cases, it is mixed with cubic diamond, graphite and other minerals. This makes it difficult ‪—‬ or even impossible ‪—‬ to test and measure its unique properties.

The new study, published March 4 in the journal Nature, addressed this problem by creating several pure hexagonal diamond samples about 0.06 inches (1.5 millimeters) in diameter ‪—‬ big enough to measure the samples' material properties. The team found that hexagonal diamond is both stiffer and harder than cubic diamond, and that it resists oxidation much more than cubic diamond does. This means hexagonal diamond can tolerate much higher temperatures without its surface getting all gunked up by reacting with oxygen, which is important for applications like drilling.

First evidence of hexagonal diamond?

The study also provides major evidence that hexagonal diamond is a real material. According to the study, "structural and spectroscopic analyses, supported by large-scale molecular dynamical simulations, unambiguously confirm the identity of HD (hexagonal diamond)."

To make the samples, the researchers compressed very organized graphite (graphite with carbon atoms neatly arranged) for 10 hours at 20 gigapascals, or about 200,000 times Earth's atmospheric pressure at sea level, and subjected them to temperatures ranging from 2,300 to 3,450 degrees Fahrenheit (1,300 to 1,900 degrees Celsius). At higher temperatures and pressures, the lonsdaleite started morphing into cubic diamond.

Hexagonal diamond could improve processes and tools that currently rely on cubic diamond, like drilling and cutting tools, polishing abrasive coatings, and dissipating heat from electronics. Its presence in meteorites can also tell us a lot about how the meteorite formed and where it came from, giving more clues about our solar system.

The elusive material "has potential applications in many fields, for example in cutting tools, in thermal management materials and in quantum sensing", Chong-Xin Shan, co-lead of the new Nature study and a physicist at Zhengzhou University, told Nature in an article.

The new study also provides "a practical strategy for producing HD (hexagonal diamond) in bulk form," opening the way for bigger samples, more scientific exploration, and industrial applications no longer limited by cubic diamond's hardness, according to the authors.

​​We tend to think of metals as hard, strong and resistant to high temperatures — just look at iron, aluminum and steel. While this is generally true, there's one key exception: mercury. With a melting point of minus 37.9 degrees Fahrenheit (minus 38.8 degrees Celsius), mercury is one of only two elements that are liquid at room temperature. (The other is bromine, which is not a metal.)

But why is mercury so different from its fellow metals?

The answer comes down to mercury's position in the periodic table and its knock-on effect on how the metal bonds together.

Melting point is directly correlated with bond strength — "the stronger the bonds, the more energy, in the form of heat, is required to break them," Zoe Ashbridge, a senior lecturer in chemistry for the U.K. Ministry of Defence, told Live Science.

Atoms of mercury, like atoms of all other metals, bind together through metallic bonding — a lattice of positively charged metal particles known as ions, is surrounded by a sea of delocalized (freed) electrons, and electrostatic attraction between these oppositely charged particles acts as the glue that holds the metal together. This structure explains many of the other signature properties of metals, such as electrical conductivity, as the electrons can move freely through the material, and mouldability, as the layers of positive particles can slide over one another to adopt a new shape, lubricated by the free electrons. But it is specifically the strength of the electrostatic attraction that governs the melting point.

The availability of outer electrons to create this delocalized sea is therefore a key factor. "The more positive the metal center is and the more delocalized valence electrons on the outside, the greater the attraction is, and generally this tracks from left to right in the periodic table," Ashbridge explained.

As a group 12 metal, mercury theoretically has 12 outer electrons it could contribute to metallic bonding. "However, all of those electrons are in "filled subshells," she added. "When they are full, that makes them more stable and less likely to delocalize, and this makes mercury particularly reluctant to share its electrons, even with other mercury atoms."

Yet this filled-subshell effect isn't big enough to explain mercury's unusually low melting point. The strength of metallic bonding — and, therefore, the melting point — also decreases from the top to the bottom of the periodic table, as the atoms get larger. But extrapolating from these established trends, mercury should still have a melting point of around 266 F (130 C), which would make it solid at room temperature.

So what causes this giant disparity?

Mercury's liquid state results almost entirely from relativistic effects, said Peter Schwerdtfeger, a quantum physicist at Massey University in New Zealand. Toward the bottom of the periodic table, the electrons in the heaviest elements experience such strong attraction to the atomic nucleus that they move close to the speed of light. At this point, they no longer obey the laws of classical physics, and the resulting quantum phenomena — known as relativistic effects — lead to surprising physical properties. How these manifest depends on the element.

"Relativistic effects become extremely important for the group 11 and group 12 elements, where gold and mercury are," he told Live Science. Consequently, the weird physical properties arising from these quantum effects are most observable in these elements. Gold has an extremely unusual yellowish hue and mercury is a liquid at room temperature.

"They show us a so-called maximum of relativistic effects, and the outer shell of these atoms contract as a result. It's enormous. For mercury, it's about 20%," Schwerdtfeger said. In chemistry terms, this relativity-induced contraction is most easily explained by once again considering mercury's electron configuration.

The full 4f subshell, which contains the electrons associated with the rare earth, or lanthanide elements, is extremely poor at shielding the other electrons from the nuclear charge. This means the outermost electrons are held much closer to the nucleus than usual — a phenomenon called lanthanide contraction. These contracted electrons move close to the speed of light and therefore experience relativistic effects.

"This increases their mass, and when they have an increased mass due to this high speed, it pulls those electrons even closer to the nucleus," Ashbridge said. Consequently, the relativistic effects reduce the availability of the electrons to contribute to metallic bonding, thus lowering the melting point of the metal below room temperature.

At a quantum mechanical level, though, this qualitative explanation is extremely challenging to back up with calculations.

"The Schrödinger equation" — which usually describes the possible positions of particles such as electrons — "doesn't fulfill the relativity principle of Albert Einstein," Schwerdtfeger explained. As a result, this equation doesn't work for high-speed particles such as the electrons in mercury. Scientists must instead turn to the significantly more complicated Dirac equation, making any simulations extremely computationally demanding.

Eventually, though, advances in computing enabled Schwerdtfeger to devise a model that could accurately simulate mercury melting and provide a quantum theoretical explanation for the anomalous melting point.

"Using what we call density functional theory, we were able to establish that the melting point is lowered by over 200 degrees Celsius [360 F] by the relativistic effects," he said. These quantum contributions dominate, so while periodic trends predict a low melting point for mercury, the relativistic effects make the element a liquid at room temperature.

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

Priestesses of a mysterious cult in ancient Greece and Rome may have used a highly toxic fungus to create psychedelic hallucinations during their rituals, a new study suggests. However, some experts say that, although the study shows that is plausible, it's not historical proof that this occurred.

The study, published Feb. 13 in the journal Scientific Reports, reports the results of laboratory experiments to make the ergot fungus non-toxic while keeping its hallucinogenic properties. A key feature of the study is that it used only the simple technology known in ancient Greece, where what's now known as the Eleusinian Mystery cult originated about 3,000 years ago.

The idea that the Eleusinian Mysteries were based on hallucinogenic substances from ergot — the "Psychedelic Eleusis" theory — has been popular since the 1970s. But the researchers are the first to show experimental evidence, Evangelos Dadiotis, a pharmaceutical scientist at the University of Athens, told Live Science in an email.

"The central question was whether toxic ergot could realistically have been processed into something psychoactive but not lethal using methods available in antiquity," he said. "We used a simple lye [sodium hydroxide] preparation made from water and ash, a technology well known in the ancient world," he said.

Wood ash produced an alkaline solution, which over time broke down the toxic proteins in the ergot (Claviceps purpurea) while leaving non-toxic byproducts, including the hallucinogenic chemical lysergic acid amide (LSA). LSA is chemically similar to lysergic acid diethylamide — better known as LSD — and it can be a precursor to the drug, but it is much less powerful.

The study suggests the ancient Greeks could have treated ergot with lye to make a non-toxic psychedelic drink for the Eleusinian Mysteries. But whether they did or not is questioned by other experts.

Mystery cult

The Eleusinian Mysteries were the most revered secret religious initiations in ancient Greece and a classic "mystery cult." They centered on the worship of the fertility goddess Demeter and her daughter Persephone, the personification of spring. According to ancient myths, the underworld god Hades abducted Persephone to be his wife and Demeter made the world barren in her grief, but the chief god Zeus made a deal so Persephone could come back each year. The cult originated in the ancient Greek town of Eleusis, hence its name. But it became popular throughout the later Roman Empire, in part because of the syncretism of Roman and Greek religious beliefs. The details of the cult are obscure ‪—‬ that was the point ‪—‬ but its initiates gathered at Eleusis each year to honor Demeter and Persephone (also called Kore) and the "mysteries" they shared, which were related to agriculture. Therefore, the mysteries were thought to have divine origins, and even Roman emperors such as Augustus became initiates, known as "mystai" in Greek.

The idea that priestesses of the Eleusinian Mysteries administered hallucinogens to initiates was proposed in the "The Road to Eleusis: Unveiling the Secret of the Mysteries" (Harcourt Brace Jovanovich, 1978) by author Gordon Wasson, classicist Carl Ruck (a co-author of the new study) and chemist Albert Hofmann, who had used an ergot derivative to make LSD in 1938 and experienced a dose himself in 1943.

But "the key objection was always toxicity ‪—‬ ergot causes ergotism, meaning convulsions, gangrene, [and] mass poisoning," Dadiotis said. Nobody had shown before that ergot could be made safe by treating it with lye, which destroyed the toxic chemicals while preserving its psychoactive properties: "Our study fills that gap … that experimental bridge is what was missing."

Strange rituals

The annual rites of the Eleusinian Mysteries were held at two times each year: the "lesser mysteries" in the spring and the "greater mysteries" in the fall. They often involved sacred processions to cult sites, ritual bathing in the sea, animal sacrifices, and fasting for several days, followed by the drinking of a mysterious elixir called kykeon, which was made with barley and flavored with herbs.

Dadiotis and his colleagues think the treated ergot extracts were added to the kykeon, and they note that in 2002 scientists said they had found traces of the psychoactive chemicals in a ceremonial vase from an Eleusinian site in Spain, and in the hardened dental plaque of an individual buried there.

The herbs added to the kykeon included a pungent type of mint, now called pennyroyal, (Mentha pulegium), and Dadiotis thinks this may have helped mask the bitter taste of the ergot extracts.

"The new study is an interesting and technically careful piece of analytical chemistry," Sharday Mosurinjohn, a religious studies scholar at Queens University in Ontario, told Live Science in an email. Mosurinjohn was not involved in the new study but has questioned the idea of the use of psychedelics in the Eleusinian Mysteries.

"What it demonstrates is chemical feasibility within a plausible ancient technological context," but "chemical feasibility is not historical proof," Mosurinjohn said. The study neither demonstrated that this type of processing was used in ancient times, nor that initiates had consumed the psychoactive doses during the Eleusinian Mystery rituals, she said.

On this day in 1940, two scientists discovered an elusive form of carbon — and inadvertently opened a window into lost civilizations.

Since the mid-1930s, scientists had predicted the existence of a form of carbon with two extra neutrons in its nucleus, but they thought it would be so short-lived that it would be impossible to measure.

But Ernest Lawrence, who founded the Berkeley Laboratory, was determined to find it. In 1939, he tasked chemists Martin Kamen and Samuel Rubin with discovering carbon-14. For a year, they found no hint of the elusive atom.

Then, in January 1940, they launched a "desperation" experiment, in which they placed a piece of graphite (a crystalline form of carbon) inside a cyclotron, one of the first types of particle accelerators. The cyclotron bombarded their sample with deuterons — nuclei of a heavy form of hydrogen with one proton and two neutrons. The hope was that the crystalline form of carbon would absorb the extra neutrons, emit a proton, and become a "heavy" version of carbon.

They ran the experiment for 120 hours straight. On Feb. 15, a sleep-deprived Kamen stopped bombarding the sample with deuterons and headed home. He was so disheveled that police, who were looking for an escaped murderer, briefly questioned him.

When Kamen was released, he returned to the lab, where his colleague Ruben noted faint signs of radioactivity in the sample. For the next two weeks, they purified the carbon, converting it into a CO2 gas that could be pumped at the right angle at the Geiger counter to measure its radioactivity.

Surprisingly, the carbon did not have a short half-life — the time it takes for half the radioactive atoms to decay into a stable atom.

"The measured cross section coupled with low yield suggests the half-life to be very long (years)," the researchers wrote in a short paper published March 15, 1940, in the journal Physical Review Letters.

Their measurements indicated it would take around 4,000 years for about half the carbon-14 to decay into nitrogen-14. (We now know the half-life of carbon-14 is about 5,730 years.)

Even at the time, they recognized the significance of their find.

"Long-lived radio-carbon will be of great importance for many chemical, biological, and industrial experiments," the researchers wrote in the paper.

In the next few years, Ruben and Kamen used radioactive carbon and oxygen molecules to elucidate the key steps in photosynthesis. Sadly, Ruben died in 1943 in a lab accident while working with a poisonous gas, and Kamen was fired from Berkeley after having social interactions with musicians and other people considered "leftists" during the Red Scare. In 1948, he was hauled up to testify in front of the House Un-American Activities Committee, and although he was never found guilty of any wrongdoing, he was dogged by unfounded allegations for years.

While the implications of Kamen and Ruben's experiments were immediately apparent, it wasn't until 1949 that University of Chicago chemists James Arnold and Willard Libby demonstrated that the ratio of carbon-14 to stable carbon could be used to estimate the ages of carbon-containing relics. Libby would earn the 1960 Nobel Prize in chemistry for his work on radiocarbon dating.

Archaeologists routinely use radiocarbon dating to estimate the ages of ancient skeletons and other artifacts that are up to 50,000 years old. And newer techniques analyze radioactive isotopes of elements such as strontium and lead to determine where ancient people lived and died, what they ate, and which pollutants they had encountered during their lifetimes.

Gleaming black teeth have been considered a high standard of beauty in parts of Vietnam since at least the late 1800s. But now, archaeologists have traced this practice back 2,000 years, discovering that ancient people used their abundant iron resources to dye their pearly whites black.

In a study published Jan. 22 in the journal Archaeological and Anthropological Sciences, researchers investigated skeletons from Dong Xa, an archaeological site in the Red River delta of northern Vietnam. The settlement at Dong Xa was occupied during the Iron Age (550 B.C. to A.D. 50), and the cemetery held numerous skeletons with unique dental colors. To figure out how people discolored their teeth thousands of years ago, the researchers nondestructively analyzed the skeletons' enamel using a variety of techniques.

When they targeted the colored areas of dental enamel using X-ray fluorescence, which measures the X-rays emitted by a sample to characterize its chemical composition, they found a high concentration of iron oxide, the researchers wrote in the study. Then, they used scanning electron microscopy with energy dispersive spectrometry (SEM-EDS). This technique involves bombarding a sample with electrons, which produces X-rays characteristic of the chemical elements in the sample. The researchers found that the discolored ancient enamel samples from Dong Xa were positive for iron (Fe) and sulfur (S).

"We believe that the combined presence of Fe and S signals is a strong indicator of the involvement of iron salts," study lead author Yue Zhang, an archaeologist at the Australian National University, told Live Science in an email. Nowadays, botanical materials are also used as part of the process to blacken teeth, so it's likely that finding traces of these on ancient teeth may also signify the practice, Zhang added.

One modern method of blackening teeth involves the combination of an iron-based substance with a tannin-rich plant material, such as betel nuts (Areca catechu). Betel-nut chewing has been popular for thousands of years among peoples of the Pacific and Southeast Asia, and prolonged use of the natural stimulant can stain a person's teeth and gums red or reddish-brown. But when tannic acids and iron salts are combined and exposed to air, they create a dark-black color.

Based on information from modern populations that blacken their teeth, the researchers suspect that the ancient blackening process likely took several days or weeks of application of an iron-tannin mixture to achieve the intensely dark shade. But once the process was completed, the person's teeth remained black throughout their lifetime, with touch-ups needed every few years to preserve their luster.

"The practice is still observed today, not only in Vietnam, but also more widely across parts of Southeast Asia," Zhang said.

While the precise procedures for tooth blackening likely changed over time, the underlying mechanism responsible for the dark coloration — the interaction between tannic acid and iron salt — was likely the same, according to the researchers. This means the presence of iron salt and sulfur on ancient teeth can be considered a diagnostic marker of purposeful blackening, they wrote.

"To our knowledge, our research on the Dong Xa teeth is the first to connect archaeologically discovered blackened teeth with modern intentional tooth-blackening practices," Zhang said.

But there are still unsolved questions surrounding why the practice of tooth blackening arose.

One possibility is that blackening was developed as a less-extreme version of tooth ablation, a practice that involves removing otherwise-healthy teeth as a rite of passage or as a group identification marker, the researchers noted. Another possibility is that blackening was invented to enhance the visual impact of the staining that resulted from betel-nut chewing.

Regardless of the original purpose, the researchers wrote, "tooth blackening plausibly became widespread around the Iron Age, when iron utensils became more accessible for producing blackening dye paste."

Human skeleton quiz: What do you know about the bones in your body?

Mineral sunscreens are notorious for leaving a white cast on the skin — but now, scientists have found that tweaking the shape of zinc oxide nanoparticles in the formula can help solve that problem.

The new, four-armed crystal structure, known as a tetrapod, was reported in December in the journal ACS Materials Letters, and it prevents individual nanoparticles from clumping together. This simultaneously boosts the stability of the sunscreen formulation and produces a warmer color that better matches a range of skin tones.

Damage caused by ultraviolet (UV) radiation in sunlight is the leading preventable cause of skin cancer in the U.S., and dermatologists strongly recommend that people apply sun protection daily to minimize their risk.

Unlike mineral sunscreens, chemical sunscreens absorb into the skin and do not leave a cast. However, many people prefer mineral sunscreens, which sit on top of the skin, but the residue leaves an unpleasant appearance after application.

"I was frustrated by how mineral sunscreen looks on my own skin," lead study author AJ Addae, a doctoral candidate in chemical biology at UCLA, said in a statement. "A lot of my motivation came from my own experience trying to use mineral sunscreen and dealing with the white cast and other unsightly aesthetic issues. This led me to simply avoid sunscreen altogether."

Mineral sunscreens use chalky zinc oxide to absorb harmful UV radiation. These products are usually formulated as colloidal materials — a suspension of fine solid particles in a watery or oily base. However, this suspended structure can create a few practical issues for both manufacturers and consumers, said Kyra Sedransk Campbell, a former professor of chemical engineering at the University of Sheffield and now CEO of Kingston Street Consulting.

For example, interactions between individual particles in the suspension form clumps over time, which can lead to visible white streaks as the product is applied to the skin. In addition, "they don't have as long a shelf life," added Sedransk Campbell, who wasn't involved in the new work. "Generally it's harder to maintain the stability of the sunscreen, and guaranteed performance is based on when the formulation is in its ideal state — so it's much harder to make the same guarantees as with a chemical sunscreen."

Addae and her colleagues decided to investigate whether changing the shape of the zinc oxide nanoparticles could help address these problems without compromising the sun protection performance.

They used a method known as "flame synthesis," which involves heating zinc with ethanol in a furnace, creating a naked flame. This creates four-armed nanoparticles, which are then incorporated into a sunscreen formulation. Scanning electron microscopy, which uses an electron beam to render high-resolution images of samples, revealed that these new particles were larger than the spherical nanostructures usually found in sunscreen but that they showed less tendency to clump together.

"Because of their structure, these tetrapod-shaped particles have standoffs and form porous networks instead of collapsing into clumps," Addae said. "They can't pack tightly and aggregate, so they stay evenly distributed in the sunscreen."

Crucially, though, this alternative shape did not affect the sunscreen's protective performance. Broad-spectrum sunscreens must protect against two wavelength bands of UV light: UVA (315 to 400 nanometers) and UVB (280 to 315 nanometers). Their trial formulation effectively absorbed both wavelength ranges and achieved a sun protection factor (SPF) of around 30, the standard for mineral sunscreens.

Product stability testing also showed that the new mixture was less prone to thickening or separating over time, compared with formulas with standard spherical particles. That means it would likely retain this high performance longer.

But a crucial element of a successful sunscreen is that people actually want to use it.

The white cast caused by clumps of zinc oxide particles in mineral formulations is a consistent complaint among U.S. consumers and cited as a reason why many people, particularly those with darker skin tones, avoid these products. The new tetrapod structures scatter visible light differently from standard spherical zinc oxide particles, thereby creating a warmer tone that would likely be more acceptable to users.

"When I spread it on my own skin, I didn't get that white cast I usually see with zinc oxide," Addae said. "That was the moment I realized this could really work."

The application-driven focus of this research is "really exciting" and has the potential to translate into real impact, Sedransk Campbell said. The next steps will involve testing the human and environmental safety profiles of the tetrapod nanoparticles and investigating how to scale up production in a cost-effective manner.

This article is for informational purposes only and is not meant to offer medical advice.

For close to a century, geoscientists have pondered a mystery: Where did Earth’s lighter elements go? Compared to amounts in the Sun and in some meteorites, Earth has less hydrogen, carbon, nitrogen and sulfur, as well as noble gases like helium — in some cases, more than 99 percent less.

Some of the disparity is explained by losses to the solar system as our planet formed. But researchers have long suspected that something else was going on too.

Recently, a team of scientists reported a possible explanation — that the elements are hiding deep in the solid inner core of Earth. At its super-high pressure — 360 gigapascals, 3.6 million times atmospheric pressure — the iron there behaves strangely, becoming an electride: a little-known form of the metal that can suck up lighter elements.

Study coauthor Duck Young Kim, a solid-state physicist at the Center for High Pressure Science & Technology Advanced Research in Shanghai, says the absorption of these light elements may have happened gradually over a couple of billion years — and may still be going on today. It would explain why the movement of seismic waves traveling through Earth suggests an inner core density that is 5 percent to 8 percent lower than expected were it metal alone.

Electrides, in more ways than one, are having their moment. Not only might they help solve a planetary mystery, they can now be made at room temperature and pressure from an array of elements. And since all electrides contain a source of reactive electrons that are easily donated to other molecules, they make ideal catalysts and other sorts of agents that help to propel challenging reactions.

One electride is already in use to catalyze the production of ammonia, a key component of fertilizer; its Japanese developers claim the process uses 20 percent less energy than traditional ammonia manufacture. Chemists, meanwhile, are discovering new electrides that could lead to cheaper and greener methods of producing pharmaceuticals.

Today’s challenge is to find more of these intriguing materials and to understand the chemical rules that govern when they form.

Electrides at high pressure

Most solids are made from ordered lattices of atoms, but electrides are different. Their lattices have little pockets where electrons sit on their own.

Normal metals have electrons that are not stuck to one atom. These are the outer, or valence, electrons that are free to move between atoms, forming what is often referred to as a delocalized “sea of electrons.” It explains why metals conduct electricity.

The outer electrons of electrides no longer orbit a particular atom either, but they can’t freely move. Instead, they become trapped at sites between atoms that are called non-nuclear attractors. This gives the materials unique properties. In the case of the iron in Earth’s core, the negative electron charges stabilize lighter elements at non-nuclear attractors that were formed at those super-high pressures, 3,000 times that at the bottom of the deepest ocean. The elements would diffuse into the metal, explaining where they disappeared to.

The first metal found to form an electride at high pressure was sodium, reported in 2009. At a pressure of 200 gigapascals (2 million times greater than atmospheric pressure) it transforms from a shiny, reflective, conducting metal into a transparent glassy, insulating material. This finding was “very weird,” says Stefano Racioppi, a computational and theoretical chemist at the University of Cambridge in the United Kingdom, who worked on sodium electrides while in the lab of Eva Zurek at the University at Buffalo in New York state. Early theories, he says, had predicted that at high pressure, sodium’s outer electrons would move even more freely between atoms.

The first sign that things were different came from predictions in the late 1990s, when scientists were using computational simulations to model solids, based on the rules of quantum theory. These rules define the energy levels that electrons can have, and hence the probable range of positions in which they are found in atoms (their atomic orbitals).

Simulating solid sodium showed that at high pressures, as the sodium atoms get squeezed closer together, so do the electrons orbiting each atom. That causes them to experience increasing repulsive forces with one another. This changes the relative energies of every electron orbiting the nucleus of each atom, Racioppi explains — leading to a reorganization of electron positions.

The result? Rather than occupying orbitals that allow them to be delocalized and move between atoms, the orbitals take on a new shape that forces electrons into the non-nuclear attractor sites. Since the electrons are stuck at these sites, the solid loses its metallic properties.

Adding to this theoretical work, Racioppi and Zurek collaborated with researchers at the University of Edinburgh to find experimental evidence for a sodium electride at extreme pressures. Squeezing crystals of sodium between two diamonds, they used X-ray diffraction to map electron density in the metal structure. This, they reported in September 2025, confirmed that electrons really were located in the predicted non-nuclear attractor sites between sodium atoms.

Just the thing for catalysts

Electrides are ideal candidates for catalysts — substances that can speed up and lower the energy needed for chemical reactions. That’s because the isolated electrons at the non-nuclear attractor sites can be donated to make and break bonds. But to be useful, they would need to function at ambient conditions.

Several such stable electrides have been discovered over the last 10 years, made from inorganic compounds or organic molecules containing metal atoms. One of the most significant, mayenite, was found by surprise in 2003 when material scientist Hideo Hosono at the Institute of Science Tokyo was investigating a type of cement.

Mayenite is a calcium aluminate oxide that forms crystals with very small pores — a few nanometers across — called cages, that contain oxygen ions. If a metal vapor of calcium or titanium is passed over it at high temperature, it removes the oxygen, leaving behind just electrons trapped at these sites — an electride.

Unlike the high-pressure metal electrides that switch from conductors to insulators, mayenite starts as an insulator. But now its trapped electrons can jump between cage sites (via a process called quantum tunnelling) — making it a conductor, albeit 100 to 1,000 times less conductive than a metal like aluminum or silver. It also becomes an excellent catalyst, able to surrender electrons to help make and break bonds in reactions.

By 2011, Hosono had begun to develop mayenite as a greener and more efficient catalyst for synthesizing ammonia. Over 170 million metric tons of ammonia, mostly for fertilizers, is produced annually via the Haber-Bosch process, in which metal oxides facilitate hydrogen and nitrogen gases reacting together at high pressure and temperature. It is an energy-intensive, expensive process — Haber-Bosch plants account for some 2 percent of the world’s energy use.

In Haber-Bosch, the catalysts bind the two gases to their surfaces and donate electrons to help break the strong triple bond that holds the two nitrogen atoms together in nitrogen gas, as well as the bonds in hydrogen gas. Because mayenite has a strong electron-donating nature, Hosono thought mayenite would be able to do it better.

In Hosono’s reaction, mayenite itself does not bind the gases but acts as a support bed for nanoparticles of a metal called ruthenium. First, the nanoparticles absorb the nitrogen and hydrogen gases. Then the mayenite donates electrons to the ruthenium. These electrons flow into the nitrogen and hydrogen molecules, making it easier to break their bonds. Ammonia thus forms at a lower temperature — 300 to 400° C — and lower pressure — 50 to 80 atmospheres— than with Haber-Bosch, which takes place at 400 to 500° C and 100 to 400 atmospheres.

In 2017, the company Tsubame BHB was formed to commercialize Hosono’s catalyst, with the first pilot plant opening in 2019, producing 20 metric tons of ammonia per year. The company has since opened a larger facility in Japan and is setting up a 20,000-ton-per year green ammonia plant in Brazil to replace some of the nation’s fossil-fuel-based fertilizer production. The company estimates that this will avoid 11,000 tons of CO₂ emissions annually — about equal to the annual emissions of 2,400 cars.

There are other applications for a mayenite catalyst, says Hosono, including a lower-energy conversion of CO₂ into useful chemicals like methane, methanol or longer-chain hydrocarbons. Other scientists have suggested that mayenite’s cage structure also makes it suitable for immobilizing radioactive isotope waste in nuclear power stations: The electrons could capture negative ions like iodine and bromide and trap them in the cages.

Mayenite has even been studied as a low-temperature propulsion system for satellites in space. When it is heated to 600° C in a vacuum, its trapped electrons blast from the cages, causing propulsion.

Organic electrides

The list of materials known to form electrides keeps growing. In 2024, a team led by chemist Fabrizio Ortu at the University of Leicester in the UK accidentally discovered another room-temperature-stable electride made from calcium ions surrounded by large organic molecules, together known as a coordination complex.

He was using a method known as mechanical chemistry — “You put something in a milling jar, you shake it really hard, and that provides the energy for the reaction,” he says. But to his surprise, electrons from the potassium he had added to his calcium complex were not donated to the calcium ion. Instead, what formed “had these electrons that were floating in the system,” he says, trapped in sites between the two metals.

Unlike mayenite, this electride is not a conductor — its trapped electrons do not jump. But they allow it to facilitate reactions that are otherwise hard to get started, by activating unreactive bonds, doing a job much like a catalyst. These are reactions that currently rely on expensive palladium catalysts.

The scientists successfully used the electride on a reaction that joins two pyridine rings — carbon rings containing a nitrogen atom. They are now examining whether the electride could assist in other common organic reactions, such as substituting a hydrogen atom on a benzene ring. These substitutions are difficult because the bond between the benzene ring carbon and its attached hydrogen is very stable.

There are still problems to sort out: Ortu’s calcium electride is too air- and water-sensitive for use in industry. He is now looking for a more stable alternative, which could prove particularly useful in the pharmaceutical industry to synthesize drug molecules, where the sorts of reactions Ortu has demonstrated are common.

Still questions at the core

There remain many unresolved mysteries about electrides, including whether Earth’s inner core definitely contains one. Kim and his collaborators used simulations of the iron lattice to find evidence for non-nuclear attractor sites, but their interpretation of the results remains “a little bit controversial,” Racioppi says.

Sodium and other metals in Group 1 and Group 2 of the periodic table of elements — such as lithium, calcium and magnesium — have loosely bound outer electrons. This helps make it easy for electrons to shift to non-nuclear attractor sites, forming electrides. But iron exerts more pulling power on its outer electrons, which sit in differently shaped orbitals. This makes the increase in electron repulsion under pressure less significant and thus the shift to electride formation difficult, Racioppi says.

Electrides are still little known and little studied, says computational materials scientist Lee Burton of Tel Aviv University. There is still no theory or model to predict when a material will become one. “Because electrides are not typical chemically, you can’t bring your chemical intuition to it,” he says.

Burton has been searching for rules that might help with predictions and has had some success finding electrides from a screen of 40,000 known materials. He is now using artificial intelligence to find more. “It’s a complex interplay between different properties that sometimes can all depend on each other,” he says. “This is where machine learning can really help.”

The key is having reliable data to train any model. Burton’s team only has actual data from the handful of electride structures experimentally confirmed so far, but they also are using the kind of modeling based on quantum theory that was carried out by Racioppi to create high-resolution simulations of electron density within materials. They are doing this for as many materials as they can; those that are confirmed by real-world experiments will be used to train an AI model to identify more materials that are likely to be electrides — ones with the discrete pockets of high electron density characteristic of trapped electron sites. “The potential,” says Burton, “is enormous.”

This article originally appeared in Knowable Magazine, a nonprofit publication dedicated to making scientific knowledge accessible to all. Sign up for Knowable Magazine’s newsletter.

The complex precursors to biological molecules can form spontaneously in interstellar space, according to a lab experiment that opens up new pathways for the origin of life in the universe.

In the presence of ionizing radiation, amino acids — the simplest units of proteins — couple together to form peptide bonds, the first step in the synthesis of more complex biological molecules such as enzymes and cell proteins, according to a new study.

The findings, published Jan. 20 in the journal Nature Astronomy, offer new possibilities for the origin of life on Earth and could help scientists target the search for extraterrestrial life, the study authors say.

The cocktail of life

Early life evolved from a complex cocktail of prebiotic molecules, including amino acids, basic sugars and RNA. But how these simple starter compounds first formed remains a mystery. One hypothesis proposes that some of these molecules may have originated in outer space and were later delivered to the early Earth through meteorite impacts, said Alfred Hopkinson, lead author of the study and a postdoctoral researcher in the Department of Physics and Astronomy at Aarhus University in Denmark.

Glycine, the simplest amino acid, is one example that has been detected in numerous comet and meteorite samples over the past 50 years, including dust samples taken from the asteroid Bennu during NASA’s recent OSIRIS-REx mission. More complex dipeptide units, which are formed when two amino acids bond by releasing water, have not been identified in these extraterrestrial bodies yet, but the intensely ionizing conditions of interstellar space gives rise to unusual chemistry and could theoretically promote the formation of these larger molecules.

"If amino acids could join in space and get to the next level of complexity [dipeptides], when that's delivered to a planetary surface, there's an even more positive starting point to form life," Hopkinson told Live Science. "It's a very exciting theory, and we wanted to see, what is the limit of complexity that these molecules could form in space?"

Remaking the universe in a lab

The team, led by Aarhus University astrophysicist Sergio Ioppolo, therefore sought to reproduce the conditions of outer space as closely as possible. Using the HUN-REN Atomki cyclotron facility in Hungary, they bombarded glycine-coated icy crystals with high-energy protons at 20 kelvins (minus 423.67 degrees Fahrenheit, or minus 253.15 degrees Celsius) and 10^-9 millibar, in order to simulate the conditions of space as closely as possible. Then, using infrared spectroscopy and mass spectrometry — methods of identifying the types of bonds present and the products’ molecular mass, respectively — the researchers analyzed the products as they formed.

Crucially, though, they used a series of deuterium labels — heavier atoms of hydrogen that produce a different signal during spectroscopic analysis — to track exactly how the glycine molecules were interacting.

Their labeled experiment quickly confirmed their initial hypothesis: The glycine molecules reacted together in the presence of radiation to form a dipeptide called glycylglycine, thus proving that more complex compounds containing peptide bonds could spontaneously form in space.

More chemical surprises

But dipeptides weren't the only complex organic molecule generated under these conditions. One surprisingly complex signal was tentatively identified as N-formylglycinamide, a subunit of one of the enzymes involved in the production of DNA building blocks and, therefore, another key player in origin-of-life chemistry.

"If you make such a vast array of different types of organic molecules, that could impact the origin of life in ways we hadn't thought of," Hopkinson said. "It's interesting to speak to other researchers — say, RNA world people — and see how that might change their picture of the early Earth."

Going forward, though, the team is investigating whether this same process occurs for other protein-forming amino acids in the interstellar medium, which would potentially open up the possibility of forming more diverse and complex peptides with contrasting chemical properties.

Scientists have developed a new technique that doubles the amount of hydrogen produced when splitting water molecules with electricity. The method works by adding a simple organic molecule and a modified catalyst to the reactor.

The adapted method lowers energy costs by up to 40% and may offer a "promising pathway for efficient and scalable hydrogen production," the researchers said in a new study published Dec. 1 in the Chemical Engineering Journal.

"Hydrogen is one of the most in demand chemicals," study co-author Hamed Heidarpour, a doctoral student at McGill University in Montreal, Canada, told Live Science. Hydrogen is used for ammonia production to produce fertilizers, in fuel cells to generate electrical energy, or burned to directly produce energy, Heidarpour said.

The main way of producing hydrogen is through steam reforming, which involves reacting water with natural gas at high temperatures and pressures to separate water's oxygen and hydrogen atoms. But these conditions mean the process is energy intensive and requires burning large amounts of fossil fuels.

Using electricity to split water into hydrogen and oxygen molecules — a method known as electrolysis — could potentially offer a way to create hydrogen with no direct carbon dioxide emissions.

This works by connecting two metal plates known as electrodes to a direct current supply and submerging the ends of the plates into water. Applying electricity to the circuit generates hydrogen at the negative electrode (anode) and oxygen at the positive one (cathode).

However, electrolysis of water is currently inefficient, expensive and uses a lot of electricity, which often comes from non-renewable sources. The main inefficiency is from producing oxygen at the anode, Heidarpour explained.

To overcome this issue, the team behind the new study adapted the standard electrolysis setup to replace the oxygen-forming reaction with one that produces hydrogen by oxidizing an organic molecule.

First, the researchers set up two chambers containing potassium hydroxide (KOH) solutions, which were separated by a thin membrane, and then connected an electrode to either chamber to form a circuit. The team added a chemical called hydroxymethylfurfural (HMF) to the anode chamber, as well as a modified copper catalyst. Heidarpour said that chromium atoms, within the surface of their specifically designed catalyst, help favor hydrogen production by stabilizing the copper atoms in their reactive state.

When the team applied electricity, electrons from the anode oxidized the aldehyde groups in the HMF molecules. This generated hydrogen and a byproduct called HMFCA, which may find use as a chemical feedstock to make bioplastics, Heidarpour said. (Aldehydes have a carbon atom doubly bonded to an oxygen atom and a single bond to a hydrogen atom.)

This adapted method effectively doubles the amount of hydrogen made in one go, when also accounting for the hydrogen created by splitting water molecules at the cathode as usual.

The reactions also ran at around 0.4 volts, which is around 1 volt lower than in conventional water electrolysis. The researchers said this helps reduce overall energy usage by up to 40%.

Heidarpour said the team is not the first to report this type of strategy but explained that they increased the overall hydrogen production rate by using a more efficient catalyst.

HMF is often made by breaking down non-food plant materials such as paper residues, making it an attractive reagent to use in these systems. However, HMF is currently an expensive material.

Other aldehyde-containing molecules such as formaldehyde could be used instead. "Where there is a surplus of low-value organic substrates, oxidizing these into more valuable chemicals with simultaneous hydrogen generation could be an attractive and environmentally-friendly way to make two feedstocks at once," Mark Symes, a professor of electrochemistry and electrochemical technology at the University of Glasgow, who was not involved in the study, told Live Science in an email.

The researchers noted that there are still ways to improve the process to make it more efficient.

For example, further work needs to be done to improve the catalyst's stability so that it "can work for thousands of hours in an industrial setting," Heidarpour said.

Scientists have converted waste cooking oil into various recyclable plastics with exceptional strength — and some were even durable enough to tow a car.

Turning nonedible waste into useful polymers is a sustainable way to create new materials, the researchers said in a new study published Nov. 28 in the Journal of the American Chemical Society.

"Waste streams offer a potentially attractive alternative to biomass-derived feedstocks [to make plastics]," the researchers wrote in the study.

One such waste stream is used cooking oil, where nearly 3.7 billion gallons is generated each year. This waste oil has so far found uses in lubricants, nonstick coatings and fuel, but much of it still gets thrown away. In the new research, the scientists found a way to convert the waste oil into useful plastic materials that are strongly adhesive and recyclable.

Oil consists of long chains of fatty acids bound to molecules of glycerol (also known as glycerin). The researchers chemically broke the oil molecules apart and then transformed the products into simpler molecules through a series of reactions.

Combining the final alcohol and ester molecules in various ways allowed the researchers to synthesize a range of polyester plastics. (Ester molecules have a carbon atom doubly bonded to an oxygen atom and also to a single oxygen atom with a carbon side chain.)

Testing the plastics' properties, including their melting points and crystallinity, revealed that these polymers are similar to low-density polyethylene (LDPE), a plastic that is commonly used in packaging and plastic bags.

The polyesters were also sticky, due to oxygen atoms in the polymer that form strong bonds with a range of materials. This is unlike LDPE, which is a hydrocarbon with only carbon and hydrogen atoms.

The researchers tested the polymer's adhesive strength by sticking two stainless-steel plates together. The plates remained tightly glued, even when up to 270 pounds (123 kilograms) of weight was attached. Towing a four-door sedan on a slightly uphill slope with these glued steel pieces also proved no trouble. This makes these polymers equal to or stronger than the commercially available adhesives that the team also tested.

These properties make these adhesives "ideal for applications in laminates and glues used in packaging, automotive components, medical devices, and electronics," the researchers wrote.

The polyester plastics were easily recycled into their original components and later remade back into plastic. Several cycles of recycling showed little impact on the plastics' properties. Some plastics could also be recycled along with other common plastics, like high-density polyethylene and polypropylene.

"This work highlights the potential of nonedible biomass waste as a renewable feedstock for…environmentally friendly alternatives to petroleum-based plastics," the researchers wrote.

Scientists have used genetically engineered bacteria to simultaneously create and color fabrics in a one-pot method. Compared with current methods that rely on fossil fuels, the new technique offers a simpler and more sustainable way to produce colored textiles.

In a new study described Nov. 12 in the journal Trends in Biotechnology, the researchers created cellulose-based fabrics spanning the colors of the rainbow by altering the conditions used to grow the bacteria.

Synthetic fibers rely "heavily on chemical synthesis and post-treatment steps that are energy-intensive, laborious, and environmentally harmful," said study lead author Sang Yup Lee, a professor in the Department of Chemical and Biomolecular Engineering at the Korea Advanced Institute of Science and Technology. Such processes can generate large amounts of greenhouse gas emissions and contaminate water and soil with heavy metals and carcinogens, Lee told Live Science in an email.

Therefore, in recent years, there has been a growing trend to use an alternative method of producing natural fibers from the fermentation of bacteria. Cellulose is a promising target, as this material mimics the natural fibers found in fabrics such as cotton. A wide range of bacteria ordinarily convert glucose into fibers of cellulose to lend structural support and defend against other microbes. However, cellulose produced by bacteria is naturally white, which means it often needs to be dyed after processing.

Lee and his team have now simplified this process by growing cellulose-producing bacteria alongside microbes that produce natural colorants. The team used strains of color-producing Escherichia coli (E. coli) to create two classes of dyes: darker violaceins (which produced colors such as purple, blue and green) and warmer carotenoids (which produced colors such as red, orange and yellow).

Initially, the researchers genetically modified the metabolic pathway of a strain of Komagataeibacter xylinus bacteria to increase cellulose production during fermentation. Subsequently adding the violacein-producing E. coli to the reaction vessel resulted in purple-, blue- and green-dyed fabric.

However, the team was not able to use the same method to achieve the warmer tones, because the bacteria did not produce enough dye to stain the cellulose fabric, likely due to poor bacterial growth. To overcome this issue, they added pregrown and treated cellulose to a culture of carotenoid-producing E. coli. This co-culture method successfully led to red-, orange- and yellow-dyed fabrics, thereby completing the team's rainbow palette.

Overall, this method "eliminates the need for separate dying and washing processes," Lee said, adding that this helps to reduce chemical waste and water consumption.

The colored bacterial cellulose showed an overall strong stability against acids, bases, heat treatments, and washing. However, the team noted that further work is needed to fully test these materials — notably, to check their durability against industrial detergents and mechanical wear and tear.

Moving forward, Lee wants to "extend the current seven color platform to a broader spectrum" and scale up the process to an industrial level while maintaining consistent quality. Further altering the way bacteria produce the cellulose could open up other uses of the material, such as biodegradable packaging, he said.

Over a feverish 10-day period in 1985, scientists conceived of a new molecule of perfect symmetry — and named it after one of the 20th century's most famous inventors and futurists.

The hunt started in the 1970s when Harry Kroto, a lab chemist at the University of Sussex in the U.K., was puzzling over the discovery of a primordial soup of organic molecules in the "vast dark clouds that lie between the stars," Kroto said in his Nobel Prize speech.

But radio and light-based data from this interstellar medium suggested there were many more long carbon-chains than should have been possible given the astrophysical molecular synthesis theories of the time. Scientists began to wonder whether cooling red giant stars were pumping the interstellar medium full of these six to eight carbon chains.

The eureka moment for Kroto was a visit to the Rice University laboratories of chemists Robert Kurl and Richard Smalley. Smalley had a special apparatus in which a laser beam vaporized atoms on the surface of a metal disk, then swept them up into a helium cloud and a vacuum to cool them, finally analyzing their makeup using another laser.

Kroto began to wonder if they could simulate the outer shells of cool red giants by swapping out the metal disk for one made of graphite, a form of carbon.

Over the first 10 days of September, the trio, along with graduate students Sean O'Brien, and Jim Heath, produced the six-to-eight-carbon chains that supported the red giant theory.

But there were some interlopers: strange forms of carbon made up of 60 carbon atoms, and a smaller concentration of an even-larger byproduct made up of 70 carbon atoms. These "uninvited guests," as Kroto called them, had actually been found in an experiment from Exxon Corporate Research Science Laboratory in New Jersey about a year earlier, but no one had paid them much attention.

After days of working, on Sept. 9, the team came to a conclusion about its structure. "C60 appeared to be really quite unreactive, а behavior difficult to reconcile with a flat hexagonal graphene sheet-the most obvious first thought," Kroto said.

In theory, a flat graphene sheet would have had tons of dangling bonds that would make it more reactive.

For many days, the scientists worked with toothpicks and jellybeans, paper cutouts of hexagons and pentagons, and other "low-tech" modeling solutions to try to puzzle out the structure of this 60-carbon molecule.

Kroto thought back to the 1967 Expo in Montreal, where 20th-century futurist and inventor Buckminster Fuller was showcasing a geodesic dome, a spherical structure with a network of triangles on its surface, which he had patented in the 1950s. Smalley went to his office to grab a book detailing Fuller’s work, and they figured out the proposed structure.

The resulting compound, which they named buckminster fullerene, was a molecule of incredible symmetry. The paper describing their new molecule was published Nov. 14, 1985 in the journal Nature, and they were soon nicknamed buckyballs.

Over the next few years, the team deduced the properties of the class of closed molecules, called fullerenes. And by 1990, scientists had figured out that by putting an electric arc between two sticks of carbon, they could produce scads of buckyballs.

Kroto, Smalley and Curl won the 1996 Nobel Prize in Chemistry for their discovery and characterization of buckyballs.

Fullerenes as a class have proven useful, and chemical relatives of buckyballs, called nanotubes, are super strong and have high thermal and electrical conductivity. These nanotubes have become crucial in atomic force microscopes, batteries, coatings and biosensors. But though scientists have proposed using buckyballs in everything from quantum computing to drug delivery, they have yet to find their niche in mainstream applications.

When you're waiting for a pot of water to heat up on the stove, tiny bubbles are the first sign it's getting ready to boil. As the water gets hotter, the bubbles get bigger, until a rolling boil signals the water has reached 212 degrees Fahrenheit (100 degrees Celsius).

Or does it? Anyone who has boiled water in a microwave will note the lack of bubbles. So, why does boiling water have bubbles, except in a microwave?

According to fluid dynamists, nanoscale bubbles constantly appear and collapse as the water heats up over a heating source like a stove. But the temperature at which noticeable bubbles start to form could sometimes be much higher than water's boiling point on paper.

"The boiling point means that at anything above that temperature, your molecules are happier being a vapor than being a liquid," said Jonathan Boreyko, a fluid dynamist at Virginia Tech. Beyond 212 F, the intrinsic energy of the water molecules — known as the chemical potential — is lower for the gas than the liquid, making the vapor the most stable form.

"But to actually execute boiling, you have to create a bubble, which has an energy cost," Boreyko told Live Science. "So just because you're happier being a vapor doesn't mean you'll successfully boil."

Therefore, the temperature at which water actually boils is a trade-off between the chemical potential energy saved by becoming a gas and the energy spent to form a bubble.

Crucially, a bubble is not just a volume of gas but also an interface between gas and liquid phases. And like all liquid interfaces, this surface is subject to surface tension.

Surface tension is a force that constantly tries to shrink the gas-liquid boundary to the smallest possible area. In the case of a bubble, this would mean collapsing entirely back into a uniform liquid. A stable bubble must therefore contain enough gas that the chemical potential energy saving is greater than the surface tension of the interface, making larger bubbles more stable.

"Surface tension is basically an energetic cost per area," Boreyko said. "Really small bubbles have a very large surface-area-to-volume ratio, whereas a bigger bubble has a smaller area relative to its volume. The volume dominates the bigger you get, which outcompetes the surface tension cost."

Consequently, water often doesn't boil until it's a little hotter than 212 F — a phenomenon known as superheating. The boiling point marks the temperature at which the gas becomes more stable than the liquid, and the extra degrees correspond to the activation energy required to create a sufficiently large bubble.

However, various factors influence how easily these bubbles can form, Mirko Gallo, a fluid dynamist at Sapienza University of Rome, told Live Science.

"Dissolved gases, impurities in the water, the surface of the container can all reduce the energy barrier for the formation of the bubble," Gallo explained. These irregularities within the bulk liquid provide a distinct nucleation point around which bubbles can form, reducing the surface tension penalty of forming a completely spherical bubble.

"If you form a bubble on an edge, it is only half a sphere, so you have a smaller surface and will need less energy," he added. "That's why the first bubbles always start appearing on the boundary of the pot."

Boiling water in microwaves

Conversely, in a microwave, the unusual heating conditions suppress bubble formation so effectively that it's possible to superheat the water by up to 36 F (20 C).

"The electromagnetic waves are penetrating and exciting the water molecules through the entire volume, so it heats the water very quickly and uniformly, whereas on a stovetop, it's the bottom wall of the pot that's getting hottest," Boreyko explained. "You also tend to [heat up things in a microwave] in a pretty smooth container — say, glass — so you don't have those localized hotspots that help you get over that energy barrier to create the first interface."

This huge store of chemical potential energy in the superheated liquid is spontaneously released in the form of a giant, explosive bubble as soon as the container is disturbed, making water heated in the microwave surprisingly dangerous.

But superheating isn't exclusive to water; it's possible for any liquid, Gallo said.

"Water has a very high surface tension compared to most liquids, but basically, the higher the surface tension, the more dramatic the effect," Boreyko added.

The frigid conditions on the surface of Saturn's largest moon, Titan, allow simple molecules in its atmosphere to break one of the most fundamental rules in chemistry, a new study shows.

According to this principle, known as "like dissolves like," mixtures containing both polar and nonpolar components, such as oil and water, usually don't mix and instead form separate layers.

But scientists from NASA's Jet Propulsion Laboratory and the Chalmers University of Technology in Sweden were surprised to discover that the polar molecule hydrogen cyanide forms stable co-crystals with the extremely nonpolar hydrocarbons methane and ethane on Titan's frigid surface — molecules that are normally entirely incompatible on Earth.

"This contradicts a rule in chemistry, 'like dissolves like,' which basically means that it should not be possible to combine these polar and nonpolar substances," lead study author Martin Rahm, an associate professor of chemistry, biochemistry and chemical engineering at the Chalmers University of Technology, said in a statement.

The new study, published July 23 in the journal PNAS, challenges a long-held pillar of chemistry and could open the door to the discovery of more exotic solid structures across the solar system.

Re-creating Titan's surface

Conditions on Titan's surface bear a striking resemblance to those of early Earth, research suggests. Its atmosphere contains high levels of nitrogen and the simple hydrocarbon compounds methane and ethane, which cycle in a localized weather system, much like Earth's water cycle.

However, until now, researchers were unsure about the fate of the hydrogen cyanide produced by reactions in this atmosphere. Is it deposited on the surface as a solid? Does it react with its surroundings? Or could it be converted into the first molecules of life?

To investigate these questions, the NASA team replicated the conditions on Titan's surface by combining mixtures of methane, ethane and hydrogen cyanide at temperatures of around minus 297 degrees Fahrenheit (minus 183 degrees Celsius). A spectroscopic analysis — a way of studying chemicals through their interactions with different wavelengths of light — yielded unexpected results, suggesting that these contrasting compounds were interacting much more closely than had ever been observed before.

It appeared that molecules of nonpolar methane and ethane had slotted into gaps in the solid crystal structure of the hydrogen cyanide — a process known as intercalation — to create an unusual co-crystal containing both sets of molecules.

Ordinarily, polar and nonpolar molecules don't mix. Polar compounds, such as water and hydrogen cyanide, have an uneven distribution of charge across the molecule, creating some areas that are slightly positive and others that are slightly negative. These oppositely charged regions are attracted to each other, forming strong intermolecular interactions between the different polar molecules and largely ignoring any nonpolar components.

Meanwhile, nonpolar oils and hydrocarbons have an entirely symmetrical arrangement of charge and interact very weakly with neighboring nonpolar molecules and not at all with polar particles. As a result, mixtures containing both polar and nonpolar components, such as oil and water, usually form distinct layers.

To explain their bizarre observations, the NASA team joined forces with researchers at the Chalmers University of Technology to model hundreds of potential co-crystal structures, assessing each for its probable stability under the conditions on Titan.

"Our calculations predicted not only that the unexpected mixtures are stable under Titan's conditions but also spectra of light that coincide well with NASA's measurements," Rahm explained.

Their theoretical analysis identified several possible stable crystal forms, which they propose are stabilized by a surprising boost in the strength of the intermolecular forces in the hydrogen cyanide solid triggered by this mixing.

Their rigorous combination of theory and experiment impressed Athena Coustenis, a planetary scientist at the Paris-Meudon Observatory in France. She is excited to see how future data, including that from NASA's Dragonfly probe (due to arrive on Titan in 2034), will complement the study's findings.

"Comparing laboratory spectra with upcoming Dragonfly mission data may reveal signatures of these solids on Titan's surface, providing insight into their geological roles and potential importance as low-temperature, prebiotic reaction environments," Coustenis told Live Science in an email. Further work could even expand this approach to other molecules likely generated by Titan’s atmosphere, including cyanoacetylene (HC₃N), acetylene (C₂H₂), hydrogen isocyanide (HNC), and nitrogen (N₂), she said. “[This] will test whether such mixing is a general feature of Titan's organic chemistry."

In 2007, scientists published a paper that laid out a recipe for a new type of biochemistry. The method would allow scientists to see what was happening in organisms in real time.

Carolyn Bertozzi, then a biochemist at the University of California, Berkeley, and her research lab had spent years trying to visualize glycans, special carbohydrate molecules that dot cell surfaces.

Glycans are one of the three major classes of biomolecules (alongside proteins and nucleic acids) and had been implicated in inflammation and disease, but scientists had found them challenging to visualize. To do so, Bertozzi built upon a chemical approach pioneered by biochemists K. Barry Sharpless, of Scripps Research, and Morten Meldal, of the University of Copenhagen.

Sharpless had laid out a vision for "click chemistry" — a way to rapidly build complex biological molecules by snapping smaller subunits together.

Biological molecules often have backbones of bonded carbon atoms, but carbon atoms aren't keen to link up. That meant that historically, chemists had to use painstaking, multistep processes that employed multiple enzymes and left unwanted byproducts. That was fine for a lab but bad for mass-producing biomolecules for pharmaceuticals.

Sharpless realized that they could simplify and scale up the process if they could snap together simple molecules that already had a complete carbon frame. They just needed a quick, powerful, reliable connector.

Separately, Sharpless and Meldal happened upon the critical connector: a chemical reaction between the compounds azide and alkyne. The trick was the addition of copper as a catalyst.

The reaction was extremely powerful and quick, and it occurred more than 99.9% of the time, without producing any byproducts.

But for Bertozzi, there was a problem: Copper is highly toxic to cells.

So Bertozzi combed the literature to devise click chemistry that was safe in living cells. She found the answer in decades' old work: Azide and alkyne would react "explosively," without the need for a catalyst, if the alkyne was forced to take on a ring shape.

In 2004, her team demonstrated that this reaction could be used to attach azide molecules to living cells without harming them. And in 2007, Bertozzi and colleagues used her method to visualize glycans within living hamster cells.

Her process involved incorporating a carbohydrate molecule modified with azide into glycans in living cells. When they added a ring-shaped alkyne molecule that was bound to a green fluorescent protein, the azide and alkyne clicked together and the glowing green protein revealed where the glycans were in the cell.

Bertozzi dubbed the process "bioorthogonal" click chemistry — so named because it would be orthogonal to — that is, would not interfere with — the biological processes occurring in the cell. Her work has proved crucial in understanding how small molecules move through living cells. It has been used to track glycans in zebrafish embryos, to see how cancer cells mark themselves safe from immune attack using the sugar molecules, and to develop radioactive "tracers" for biomedical imaging. And click chemistry more broadly has supercharged the process of drug discovery.

In 2022, Sharpless, Meldal and Bertozzi earned the Nobel Prize in chemistry for their work on click chemistry.

Three scientists who developed materials that are like something out of "Harry Potter" have won the 2025 Nobel Prize in Chemistry.

Susumu Kitagawa of Kyoto University in Japan, Richard Robson of the University of Melbourne in Australia and Omar M. Yaghi of the University of California, Berkeley, won the prestigious prize "for the development of metal-organic frameworks."

The frameworks included new materials that store large amounts of gas in a tiny volume, "almost like Hermione's handbag in Harry Potter," Heiner Linke, chair of the Nobel Committee for Chemistry, said in a news conference announcing the prize.

In the Harry Potter series, Hermione Granger has a magical handbag that allows her to carry more objects than it should be possible to fit — like the handbag equivalent of the TARDIS from Doctor Who.

Unlike Hermione's bag, the new molecular architecture developed by Nobel laureates doesn't break the laws of physics, but it works in a similar way. The metal-organic frameworks are made up of metal ions linked together by long organic molecules. These ions and molecules are arranged in such a way that they form crystals with large cavities, which can be used to capture and store substances.

The Royal Swedish Academy of Sciences announced the winners at a ceremony in Stockholm, Sweden, on Wednesday (Oct. 8). This is the 117th Nobel chemistry prize and comes with a cash prize of 11 million Swedish kronor ($1.2 million).

"I'm deeply honored that my long-standing research has been recognized," Kitagawa said by phone at the news conference, adding that the most significant potential application of his work is in separating materials from the air "which contain most of the elements for our important materials."

The researchers' work began in 1989, when Robson combined positively-charged copper atoms to a four-armed molecule to form a spacious crystal, like a diamond filled with countless tiny compartments. Yaghi and Kitagawa followed up on this between 1992 and 2003 with work that showed gases could flow in and out of this metal-organic framework; while also making it more stable, flexible and modifiable for unique properties.

Since these discoveries, other researchers have developed innumerable metal-organic frameworks, using them to capture toxic gases required to make semiconductors, harvest water from the desert air, catalyze chemical reactions, and break down harmful chemicals and pollutants — including PFA "forever chemical" plastics, pharmaceutical runoff and chemical weapons.

But the biggest application of the frameworks could lie in the future. They are currently being tested in capturing the carbon dioxide released by factories and power stations.

"My dream is to capture air and separate air," Kitagawa said. "For instance in CO₂, or oxygen, or water, and convert this to useful materials using renewable energy."

The second you cut open an avocado, the clock starts ticking down. Within hours, the fruit goes from an appetizing green to an unappealing brown. But why do avocados turn brown so quickly?

It comes down to chemistry. Avocados have an enzyme called polyphenol oxidase (PPO).

"When you cut open an avocado, you are exposing the enzyme in the avocado cells to oxygen in the air," Matthew Fatino, the subtropical crops adviser at the Cooperative Extension of University of California Agriculture and Natural Resources, told Live Science. PPO catalyzes the reaction of the avocado's phenolic compounds, a large class of small molecules that have antioxidant and aromatic properties, with oxygen. This oxidation produces a pigmented compound called melanin, which is brown in avocados.

The oxidation process, often called enzymatic browning, breaks down the fruit. "Avocados are superhigh in fat content," Fatino said. "Humans crave that nice, buttery, fat texture." So, as the enzymes react with oxygen, the creamy fat in the fruit can start turning bitter, which isn't so appealing.

A bit of browning doesn't mean you should throw the avocado away, and you can still eat it.

"Only the exposed flesh will turn brown, and you can remove that thin layer and enjoy the green part underneath," said Sarah Alsing, a registered dietitian and recipe creator for Delightfully Fueled.

Mashing up the avocado can also hide the slight bitterness. But there's a limit, Fatino said. "If you let it go too far — I'm talking like days — you can kind of get a rancid taste," he said.

How to make avocados last longer

You may have heard this trick to delay the browning process: Keep the pit in, if you're not ready to use the whole avocado.

"It's because the pit is covering a lot of the cells," Fatino said. Some browning might occur around the pit, but it'll be green below the pit.

The news isn't so good for your leftover guac. "Conversely, if you smash up an avocado, it'll brown way more quickly, because there is more surface area exposed to oxygen," he added. However, covering unused avocado or prepared guacamole with plastic wrap or putting it in an airtight container limits the oxygen exposure and can help to preserve it.

Another way to keep an avocado fresh is to lower its pH so it's more acidic. That's why cut-up, packaged fruit often has lemon or lime juice in the ingredient list as a preservative — and it could help to add a little extra lime to your guacamole.

"The citric acid in lemon and lime juice delays oxidation," Alsing said. "Squeeze a little juice over the avocado flesh to delay browning."

In addition to turning brown, avocados sometimes have brown strands running through them.

"The fibrous threads in avocados are called vascular bundles, and they carry nutrients and water to help the avocado grow," Alsing told Live Science. "These are usually only seen in avocados from immature trees."

Avocados can also become fibrous if they overripen on the tree. If the fruits stay on too long past the season, the avocado pit — the fruit's seed — starts to germinate. "The seed will probably pull more photoassimilates [carbohydrates created during photosynthesis] from the tree, and the vascular bundles are going to further develop to nurse that seed," Fatino said.

Environmental stressors can also affect how long avocados last. The avocado tree — which is native to Mexico and Central America — is sensitive to frost and extreme heat.

"As little as a couple of degrees colder is enough to damage avocado buds and avocado trees to where it could ruin your crop for the next year," Fatino said. Too much heat makes the tree shut down and drop fruit.

"With any kind of physical damage, oxygen can get in and the breakdown of the fruit is going to occur faster," Fatino said.

So whether you're mashing it for guacamole, slicing it over toast, or scooping it to eat plain, don't sweat a bit of browning on an avocado. Enjoy the tasty fruit, keep oxygen away from any leftover avocado, and squeeze on a little citrus juice.

Scientists have discovered that the molecular building blocks needed for life are "readily available" on Saturn's icy moon Enceladus.

At only 314 miles (505 kilometers) wide, Enceladus could fit inside Colorado — and thanks to its liquid water, hydrothermal energy source and chemical tool kit, it has the potential to host extraterrestrial life.

Twenty years ago, NASA's Cassini spacecraft discovered evidence that a vast salty ocean hidden beneath Enceladus' surface was spitting out minuscule "ice grains" through cracks near the moon's south pole. Subsequent studies have spotted five of the six essential elements for life — carbon, hydrogen, nitrogen, oxygen and phosphorus (only missing sulfur) — within these grains.

However, the majority of these past studies looked at the relatively old ice grains that settled in Saturn's E ring — a diffuse ring outside the planet's bright main rings — after being ejected decades or centuries prior. This meant scientists couldn't be sure that the compounds truly came from Enceladus rather than from space weathering in the ring.

Now, astronomers have identified organic molecules, perhaps including nitrogen and oxygen, in fresh ice grains sprayed from Saturn's icy moon. The new research was published Wednesday (Oct. 1) in the journal Nature Astronomy.

Secrets of the ice moon

In 2008, as Cassini shot through a geyser of freshly spewed-up ice grains from Enceladus, it collected data on the splatter that covered the spacecraft's Cosmic Dust Analyzer. These grains hit the spacecraft at 11 miles per second (18 kilometers per second), which was so fast that the water molecules didn't cluster. This meant the team could see "previously hidden signals," study co-author Nozair Khawaja, a planetary scientist at the Free University of Berlin, said in a statement.

The researchers used mass spectrometry to analyze the chemical fingerprint of the molecules in the fresh ice grains. They found chemical compounds that, on Earth, are involved in reactions that lead to the formation of complex molecules required for life, including structures potentially containing nitrogen and oxygen.

"These molecules we found in the freshly ejected material prove that the complex organic molecules Cassini detected in Saturn's E ring are not just a product of long exposure to space, but are readily available in Enceladus's ocean," study co-author Frank Postberg, a professor of planetary science at the Free University of Berlin, said in the statement.

Nozair said there are various ways these molecules could become biologically relevant, "which enhances the likelihood that the moon is habitable." Even so, he said it would still be a huge discovery to not find any life on Enceladus because it would raise "serious questions about why life is not present in such an environment when the right conditions are there."

ESA is planning a future mission to land a spacecraft on the southern pole of Enceladus to collect more samples. The agency is targeting the early 2040s as the earliest possible launch date.

Scientists have developed a new method to break down harmful "forever chemicals" by exposing them to a sunlight-activated material.

Perfluoroalkyl and polyfluoroalkyl substances (PFAS) are chemicals found in many household products, including cookware, cosmetics, dental floss and waterproof clothing. True to their nickname, the chemicals take thousands of years to break down, enabling them to accumulate in the environment and our bodies.

PFAS have been used since the 1940s. Initially, they were valued for their nonstick properties, but now they are linked to a number of health impacts, including increased risks of autoimmune disease, developmental disorders, reduced fertility and cancer in humans. This has led some PFAS to be banned. But with nearly 15,000 types having been produced, roughly 98% of the U.S. population has these chemicals in their blood.

Now, a team of researchers has found a way to break down the chemicals, reducing them to components that include fluoride, which is harmless at low doses. They published their findings July 25 in the journal Small.

"PFAS contamination continues to pose a global health risk, and this research represents a critical step toward safer communities and cleaner ecosystems," lead researcher Cameron Shearer, a materials scientist at the University of Adelaide in Australia, said in a statement.

Related: Scientists find a simple way to destroy 'forever chemicals' — by beheading them

PFAS owe their persistence to their strong chemical bonds; they consist of a head (often charged oxygen molecules) linked to a tail of carbon and fluorine atoms. For PFAS to degrade, this bond must be broken — but this process is very difficult to achieve using traditional methods.

"Many water contaminants are degraded by adding a reactive chemical that binds to the carbon," Shearer said. "However, in PFAS molecules, the carbon atoms are protected in such a way that makes this process nearly impossible."

In recent years, researchers have been developing methods to break down PFAS using materials called photocatalysts, which absorb incident light to speed up chemical reactions. The scientists behind the new study turned to a photocatalytic material called cadmium indium sulfide, known for its ability to release reactive oxygen species — or free radicals — after being exposed to visible light.

After mixing the material with one common PFAS called perfluorooctane sulfonate (PFOS), the researchers watched as the photocatalyst absorbed light to generate free radicals that attacked the fluorine atoms in the bond.

Under optimized conditions, this led to the "complete breakdown" of around 99% of the PFOS molecules. The byproducts were components which the scientists say can be isolated and used to make toothpaste and fertilizer additives.

"The materials we have developed through our research could be used as part of PFAS-treatment chains that first capture and concentrate PFAS in water, which can then be degraded through exposure to our light-activated materials," Shearer said. "We plan to build on this study through our ongoing work improving the stability of the materials before they can be applied to large scale systems."

Scientists have created the first sizable meteorite diamond — also known as lonsdaleite or hexagonal diamond — a material predicted to be even harder than the diamonds normally found on Earth.

The high-pressure, high-temperature technique created tiny disks of this ultrahard diamond which could ultimately replace conventional diamonds in applications such as drilling tools and electronics, the scientists reported July 30 in the journal Nature.

Diamonds hold the record for the world's hardest naturally occurring substance. Each carbon atom in the infinitely repeating molecular structure forms four equal-length bonds to other carbon atoms, each separated by a 109.5 degree angle, to create an endless array of perfect tetrahedra. Viewed from the side, this structure appears to contain three repeating layers of carbon atoms (labeled A, B, and C), and this gives diamond what crystallographers call a face-centered cubic crystal structure.

In the 1960s, however, a subtly different structure of diamond was proposed, with small impure crystals of this structure subsequently discovered in the Canyon Diablo meteorite, which crashed in the Arizona desert around 50,000 years ago.

Unlike in cubic diamond, this form contains two different bond lengths — one slightly longer than in normal diamond and one slightly shorter. The carbon atoms are still organized into endless planes of tetrahedra. But this time, when viewed from the side, the structure contains only two repeating layers (labeled A and B). This slight shift in the carbon layers gives meteorite diamond a hexagonal structure, which scientists theorize should boost the solid's hardness by 58%.

But preparing samples of this hexagonal structure large enough to analyze has been challenging. What's more, the presence of other contaminating forms of carbon in the original meteorite sample — including graphite, cubic diamond and amorphous carbon — led many to doubt whether hexagonal diamond exists at all.

Related: Why do diamonds come in different colors?

Inspired by the Canyon Diablo meteoric fragment, Wenge Yang and colleagues at the Center for High Pressure Science and Technology Advanced Research in Beijing, sought to reproduce the intense conditions of an impact with Earth in the lab, developing a high-pressure, high-temperature synthesis using a diamond anvil cell, a piece of equipment that squashes a sample between two flattened surfaces made of diamond. Starting from another form of carbon, purified graphite, they slowly and carefully compressed the material, fixing the shifted atoms in place with targeted heat from a laser.

"At pressures around 20 GPa (200,000 atmospheres), the flat carbon layers of graphite are forced to slide and bond with adjacent layers, forming a buckled carbon honeycomb characteristic of hexagonal diamond," Yang told Live Science in an email. "Laser heating above 1400 °C [2,552 Fahrenheit] facilitates this transition." Once these distorted tetrahedra of hexagonal diamond had formed, the team slowly released the pressure, ensuring the new crystal didn't spontaneously turn back into graphite.

The team then used powerful techniques to view the crystal structure and confirm their achievement. Although the crystal disk remained somewhat impure, containing random fragments of cubic diamond, electron microscope images clearly showed its AB carbon layers, and X-ray crystallography revealed the hexagonal structure.

"It's a good first demonstration," said Soumen Mandal, a physicist who specializes in the applications of diamond at the University of Cardiff in the U.K., who was not involved in the study. "Now we need pure crystals and more material to start exploring its physical and mechanical properties, thermal properties, electric properties, all of these."

Hardness testing generally requires larger samples than the ones Yang's team produced, according to the study. However, they did confirm the new material was at least as tough as regular diamondsl and Yang hopes subsequent experiments with larger and purer crystals will soon provide a concrete answer.

The team would ultimately like to see hexagonal diamond begin to replace conventional diamond in industrial technologies such as precision machinery, high-performance electronics, quantum technologies and thermal management systems, although such applications may still be 10 years away.

"Looking forward, our goal is to produce larger, high-quality hexagonal diamond samples suitable for real-world applications," he said. "These efforts will help tailor hexagonal diamond's properties for specific applications and pave the way for its industrial adoption."

Chemists have synthesized a "super alcohol" previously thought to be too unstable to exist. The molecule was formed under extreme, "space-like" conditions and may shed light on the complex reactions required to form extraterrestrial life.

The super alcohol is called methanetetrol and is the only alcohol with four oxygen and hydrogen groups around a single carbon atom. It is thought to be a key building block for life in the universe. "This work pushes the boundaries of what we know about chemistry in space," study co-author Ralf Kaiser, a chemist at the University of Hawaii, said in a statement.

Methanetetrol is a type of ortho acid, a class of compounds thought to be critical in the chemistry of early life. However, these compounds are hard to isolate and study. Methanetetrol's high number of oxygen bonds, for example, make it very unstable and likely to break down if not kept in specific environmental conditions.

Methanetetrol doesn't occur naturally on Earth, but scientists have theorized about its existence and chemical structure for more than a century, according to the statement.

To emulate how methanetetrol might form in space, the researchers put water and carbon dioxide in a minus 451 degree Fahrenheit (minus 268 degrees Celsius) cryocooler, and then exposed the mixture to cosmic ray-like radiation to trigger the necessary chemical reactions to bring these molecules together.

Using ultraviolet light, the team then detected small amounts of the super alcohol in a gaseous form. They revealed their findings in a study published July 14 in the journal Nature Communications.

Related: Why is alcohol used to preserve things?

This work "pushes experimental and detection capabilities to the 'final frontier, the next level beyond what could be accomplished before due to the lack of experimental and computational approaches," Kaiser said in another statement.

"A prebiotic bomb"

Now that they are able to study methanetetrol, astrobiologists can learn how this alcohol is synthesized and reacts with other molecules in the coldest parts of space, such as the dust clouds where stars and planets form.

"You have this compact, carbon-oxygen molecule that just really wants to go 'boom,'" study co-author Ryan Fortenberry, an astrochemist at the University of Mississippi, said in the statement. "And when it does, when you give it any kind of energy, you'll have water, hydrogen peroxide and a number of other potential compounds that are important for life."

Methanetetrol is like "a seed of life molecule," he explained, that could be a building block of life across the universe. "It's something that can lead to more complex chemistry if given the opportunity."

Fortenberry likened the alcohol to an acorn, which can't grow into a tree on its own, but requires the right conditions and reactions to do so. "It's like a prebiotic bomb," he said.

The chemists concluded that, because it can form in the space-like conditions of a lab, methanetetrol can also form in space. "If we can find places where methanetetrol forms naturally, we know that it is a place that has the potential building blocks to support life," Fortenberry said.

Scientists have used ultrafast high-intensity lasers to superheat gold to 14 times its melting point without turning the solid metal into a liquid.

The record-breaking experiment, which was described in a study published July 23 in the journal Nature, smashed a decades-old theory about the stability of solids and is the first reliable method to precisely measure the temperature of extremely hot systems, the researchers said.

Unusual states of matter, such as the plasma surrounding the sun or the high-pressure cores of planets, can reach incredible temperatures of millions of degrees Fahrenheit. However, actually putting a figure to this so-called "warm dense matter" has proven challenging, as scientists have struggled to measure the short-lived hot material fast enough to get reliable results.

"We have good techniques for measuring density and pressure of these systems, but not temperature," study co-lead author Bob Nagler, a scientist at the Department of Energy's SLAC National Accelerator Laboratory, said in a statement. "In these studies, the temperatures are always estimates with huge error bars, which really holds up our theoretical models — it's been a decades-long problem."

Speed was therefore key to taking a successful measurement. To achieve this, the team used 45-femtosecond (45 quadrillionths of a second) X-ray laser pulses to rapidly heat a thin gold film. As the radiation passed through the crystalline film, the atoms vibrated at a frequency directly related to their increasing temperature. A second pulse fired at the hot sample then scattered off these vibrating atoms, and the shift in frequency of these deflected beams provided a quantitative measurement of the atoms' speed and therefore temperature.

However, the researchers realized that they had achieved much more than a new measurement technique. "We were surprised to find a much higher temperature in these superheated solids than we initially expected, which disproves a long-standing theory from the 1980s," study co-lead author Thomas White, an associate professor of physics at the University of Nevada, Reno said in the statement.

Related: Scientists spot water molecules flipping before they split, and it could help them produce cheaper hydrogen fuel

The solid gold sample reached a staggering 19,000 kelvins (33,700 degrees Fahrenheit, or 18,700 degrees Celsius) — 14 times the element's standard melting point of 1,337 kelvins (1,947 F, or 1,064 C). "This is possibly the hottest crystalline material ever recorded," White added in another statement. "I was expecting the gold to heat quite significantly before melting, but I wasn't expecting a fourteen-fold temperature increase!"

Normally, solids and liquids have a defined temperature at which they change from one state to another. But under certain conditions, materials can be heated beyond these limits without changing state — a phenomenon known as superheating. This effect is sometimes seen in water heated in a microwave. If the container is smooth, there are no irregularities around which bubbles can form so the liquid water bypasses 212 F (100 C) without boiling. However, the slightest disturbance can trigger "catastrophe," and the water explosively boils as this metastable state is broken.

In the 1980s, physicists calculated the limit of this superheating effect for solids as three times the melting point, which they dubbed the "entropy catastrophe." Above this point, the solid would theoretically have a greater entropy, or disorder, than its liquid form, breaking the second law of thermodynamics. As this law states that entropy must always increase, the idea that the neatly arranged particles of a solid could be more disordered than the random distribution of particles in a liquid is an impossible contradiction.

So how did the gold sample remain solid at 14 times its melting point? The team suggested that the sheer speed at which they heated the gold prevented the crystal structure from expanding during the timescale of the experiment.

"It's important to clarify that we did not violate the second law of thermodynamics," White said. "What we demonstrated is that these catastrophes can be avoided if materials are heated extremely quickly — in our case, within trillionths of a second."

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

For the first time, researchers have recreated the universe's first ever molecules by mimicking the conditions of the early universe.

The findings shake up our understanding of the origin of stars in the early universe and "calls for a reassessment of the helium chemistry in the early universe," the researchers wrote in the new study, published July 24 in the journal Astronomy and Astrophysics.

The first stars in the universe

Just after the Big Bang 13.8 billion years ago, the universe was subject to extremely high temperatures. A few seconds later, though, temperatures decreased enough for hydrogen and helium to form as the first ever elements. Hundreds of thousands of years after those elements formed, temperatures became cool enough for their atoms to combine with electrons in a variety of different configurations, forging molecules.

According to the researchers, a helium hydride ion — or HeH+ — became the first ever molecule. The ion is needed to form molecular hydrogen, now the most abundant molecule in the universe.

Both helium hydride ions and molecular hydrogen were critical to the development of the first stars hundreds of millions of years later, the researchers said.

For a protostar to begin fusion — the process that enables stars to create their own energy — atoms and molecules within it must collide with each other and release heat. This process is largely ineffective at temperatures under 18,000 degrees Fahrenheit (10,000 degrees Celsius).

Related: Universe's First Molecule Detected in Space for the First Time Ever

However, helium hydride ions are particularly good at continuing the process, even under cool temperatures, and are considered to be a potentially integral factor of star formation in the early universe.

The amount of helium hydride ions in the universe may therefore have had significant bearing on the speed and efficacy of early star formation, the researchers said in a statement.

Far more important than previously assumed

In the new study, the researchers recreated early helium hydride reactions by storing the ions at minus 449 degrees Fahrenheit (minus 267 degrees Celsius) for up to 60 seconds to cool them down before forcing them to collide with heavy hydrogen. Researchers studied how the collisions — similar to those that kickstart fusion in a star — changed depending on the temperature of the particles.

They found that reaction rates between these particles do not slow down at lower temperatures, which contradicts older assumptions.

"Previous theories predicted a significant decrease in the reaction probability at low temperatures, but we were unable to verify this in either the experiment or new theoretical calculations," study co-author Holger Kreckel, who studies nuclear physics at the Max Planck Institute for Nuclear Physics in Germany, said in the statement.

This new finding of how helium hydride ions function challenges how physicists think stars formed in the early universe. Reactions between the ions and other atoms "appear to have been far more important for chemistry in the early universe than previously assumed," Kreckel said.

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

Globs of sticky goo discovered in the bottom of 2,500-year-old bronze jars from southern Italy have been chemically identified, settling a 70-year archaeological debate.

It's honey — the sweet leftovers of an offering to an ancient god.

A team of chemists and archaeologists used cutting-edge analysis techniques to test the paste-like residue. They concluded that the jars, which were found in the sixth-century-B.C. city of Paestum, originally contained honeycomb.

"What I find interesting is that the ancient Greeks did think that honey was a superfood," study lead author Luciana da Costa Carvalho, a chemist at the University of Oxford, said in a video. The researchers published their findings Wednesday (July 30) in the Journal of the American Chemical Society.

Honey and honeybees were important in ancient Greek and Roman medicine, rituals, cosmetics and food. So when archaeologists found eight bronze jars in an underground shrine in 1954, they assumed that the jars contained honey as a symbol of immortality. Despite at least four attempts over seven decades to confirm the presence of the sticky, sweet substance, no evidence of sugars was ever found.

Related: Does honey ever go bad?

But Carvalho and colleagues decided to take advantage of recent advances in chemical analysis techniques and to reopen the question of the gooey substance's origin.

Using mass spectrometry, a technique that can identify different molecules and compounds, Carvalho and colleagues identified intact hexose sugars in the ancient jar residue for the first time. Fresh honey is about 79% hexose sugars, the researchers wrote in the study, with fructose being the most abundant.

An analysis of the proteins in the ancient sample revealed the presence of royal jelly, a milky secretion made by worker bees. The researchers also recovered peptides — short amino acid chains that are smaller versions of proteins — unique to one species of honeybee: the European honeybee (Apis mellifera).

Adding up these analyses, the researchers wrote that the study presents the first direct molecular evidence supporting the presence of honey, likely offered as honeycombs.

"The amount of sugar in the ancient residue is very low compared to modern honey," Carvalho told Live Science in an email. "I think the residue tastes like washed honeycomb but slightly more acidic," Carvalho said, although she did not actually try it.

The researchers also identified copper ions in the honey mixture. Because these ions are biocidal, meaning they can kill microorganisms, "their presence would have contributed to the preservations of sugars on the surface of the residue," Carvalho said, potentially explaining how the honey lasted thousands of years.

The analysis of the goop can help archaeologists better understand ancient rituals and shrines. The jars were found in an underground shrine, also called a heroon, at Paestum. The heroon also included a large, wooden table with wool-wrapped iron rods placed on top.

The offering may have been made to Is of Helice, the mythical founder of the ancient Greek city of Sybaris, located in what today is the arch of Italy's boot. When Sybaris was destroyed in the sixth century B.C., its inhabitants fled and founded a city called Poseidonia. But when the Romans took it over in the third century B.C., they renamed the city Paestum.

The new study shows that "there is merit in reanalyzing museum collections because analytical techniques continue to develop," Carvalho said in the video.

​​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?

These mysterious "fireworks" aren't lighting up the night sky — they're computer simulations from a recent paper on mixing fluids that don't want to mix.

Researchers mapped out how two immiscible fluids (two fluids that do not mix, like oil and water) with different viscosities can create "fingers" when they interact. They created different patterns by alternately injecting the fluids at the center of each "firework," allowing the fluids to spread out.

Studying this phenomenon is important for storing carbon from the atmosphere in the ground, a strategy for tackling climate change. Carbon dioxide is responsible for about 80% of all heating from human-caused greenhouse gases since 1990. Removing large amounts of carbon dioxide from the atmosphere is possible, but it still has to go somewhere. Storing it in the ground is one option — and understanding fluid interactions can help us figure out how to do that.

In this case, the word "fluid" can refer to both gases and liquids, including gaseous carbon dioxide. Viscosity is a measure of how easily a fluid moves. Fluids with high viscosity move sluggishly, like molasses or tar, while low-viscosity fluids move faster and can spread out more, like water or air.

The fluid "fireworks" are caused by Saffman-Taylor instability — a phenomenon that occurs when two immiscible fluids with different viscosities are confined in a small space. When a less viscous fluid is added to the system, there aren't a lot of places for it to go, so it pushes against the thicker fluid instead — forming the distinctive patterns.

If you've ever put a drop of glue between two flat surfaces, then changed your mind and pulled them apart, you might have noticed the wet glue forming strange ridges and channels. This is Saffman-Taylor instability in action. When you pulled the pieces apart, air tried to go where the more viscous glue was and left those patterns behind.

Related: Why does nearly all life breathe oxygen?

Storing carbon dioxide in the ground involves "injecting" carbon dioxide gas into a more viscous liquid (water) in confined spaces underground, leading to Saffman-Taylor instability. The "fireworks" from the paper show that the number and extent of the fingers can be changed depending on when and how the fluid is injected into the system. Increasing the fingering effect helps keep the gas from escaping back into the atmosphere.

People across the globe are already working on carbon sequestration (storage) projects — as of 2024, there were 50 facilities in operation, 44 being built, and an additional 534 in development according to the Global CCS (carbon capture and storage) Institute. Developing this technology further gives us more tools to rein in global heating caused by the presence of too much carbon dioxide in Earth's atmosphere.

Do you think you've got decent science knowledge? It's time to put your gray matter to the test with our weekly, free science crossword puzzle.

We've spent hours carefully writing our puzzles to make sure they are challenging but accessible to people of all ages and scientific backgrounds, with answers ranging from unusual animals and ancient rulers, to fundamental theories and Nobel Prize-winning scientists. Don't expect it to be easy, but it will be fun!

All you have to do to play is register once and then you should be logged in for next time, and if you need a hint tap the question mark next to the clue to reveal a letter.

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Want to try your luck with our previous crossword puzzles? The most recent ones can be found below, but you can access the full list of science crosswords here.

—#53: Ancient supercontinent — 11 across

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—#50: Longest-serving president in US history — 1 across

—#49: 'Short' tempered French emperor — 13 across

In a first, researchers have developed a new copper alloy that's one of the most resilient copper-based materials ever made.

The new alloy, a mixture of copper, tantalum and lithium, was built on nanoscales to withstand extreme temperatures and strains, and could have crucial applications for aerospace, defense and industry. The researchers published their findings March 27 in the journal Science.

"This is cutting-edge science, developing a new material that uniquely combines copper's excellent conductivity with strength and durability on the scale of nickel-based superalloys," study co-author Martin Harmer, an professor emeritus of engineering at Lehigh University in Bethlehem, Pennsylvania, said in a statement.

Currently, the most common materials used in high-stress environments such as gas turbine engines and chemical processing equipment are nickel-based superalloys, which are strong, resistant to corrosion and can withstand high temperatures.

Related: Scientists discover revolutionary method that makes fuel from water and sunlight — but it's not finished yet

But these alloys fall short in terms of their electrical conductivity, limiting some of their potential applications. To solve this problem, the researchers sandwiched copper-lithium precipitates between two layers rich in tantalum, an element that is highly resistant to corrosion.

The team then refined the substance further by adding a tiny amount of lithium to change the precipitates' structure into stable cuboids, bolstering the alloy's strength and thermal resilience

"When we look inside our body, we try to look for fingerprints of cell mutation for cancer," study co-author Kiran Solanki, a professor of engineering at Arizona State University, said in a statement. "Similarly, structural materials have a unique fingerprint when they are subjected to any event like radiation or heat. And in this case, having a copper lithium precipitate with a stable bilayer of Ta [tantalum] is when we can alter high temperature fingerprint for failure."

The resulting material has an impressive combination of properties. Alongside its electrical conductivity, it can operate at temperatures up to 1,472 degrees Fahrenheit (800 degrees Celsius) and can withstand a maximum stress of 1,120 megapascals at room temperature — more than one and a half times the maximum pressure that steel can endure.

These traits mean it could be used in a variety of ways, the researchers said.

"It provides industry and the military with the foundation to create new materials for hypersonics and high performance turbine engines," Harmer said.

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

Nearly three fourths of Earth is covered by oceans, making the planet look like a pale blue dot from space. But Japanese researchers have made a compelling case that Earth's oceans were once green, in a study published in Nature.

The reason Earth's oceans may have looked different in the ancient past is to do with their chemistry and the evolution of photosynthesis. As a geology undergraduate student, I was taught about the importance of a type of rock deposit known as the banded iron formation in recording the planet's history.

Banded iron formations were deposited in the Archean and Paleoproterozoic eons, roughly between 3.8 and 1.8 billion years ago. Life back then was confined to one cell organisms in the oceans. The continents were a barren landscape of grey, brown and black rocks and sediments.

Rain falling on continental rocks dissolved iron which was then carried to the oceans by rivers. Other sources of iron were volcanoes on the ocean floor. This iron will become important later.

The Archaean eon was a time when Earth's atmosphere and ocean were devoid of gaseous oxygen, but also when the first organisms to generate energy from sunlight evolved. These organisms used anaerobic photosynthesis, meaning they can do photosynthesis in the absence of oxygen.

It triggered important changes as a byproduct of anaerobic photosynthesis is oxygen gas. Oxygen gas bound to iron in seawater. Oxygen only existed as a gas in the atmosphere once the seawater iron could neutralize no more oxygen.

Eventually, early photosynthesis led to the "great oxidation event", a major ecological turning point that made complex life on Earth possible. It marked the transition from a largely oxygen free Earth to one with large amounts of oxygen in the ocean and atmosphere.

The "bands" of different colors in banded iron formations record this shift with an alternation between deposits of iron deposited in the absence of oxygen and red oxidized iron.

The case for green oceans

The recent paper's case for green oceans in the Archaean eon starts with an observation: waters around the Japanese volcanic island of Iwo Jima have a greenish hue linked to a form of oxidized iron - Fe(III). Blue-green algae thrive in the green waters surrounding the island.

Despite their name, blue-green algae are primitive bacteria and not true algae. In the Archaean eon, the ancestors of modern blue-green algae evolved alongside other bacteria that use ferrous iron instead of water as the source of electrons for photosynthesis. This points to high levels of iron in the ocean.

Photosynthetic organisms use pigments (mostly chlorophyll) in their cells to transform CO₂ into sugars using the energy of the sun. Chlorophyll gives plants their green color. Blue-green algae are peculiar because they carry the common chlorophyll pigment, but also a second pigment called phycoerythrobilin (PEB).

In their paper, the researchers found that genetically engineered modern blue-green algae with PEB grow better in green waters. Although chlorophyll is great for photosynthesis in the spectra of light visible to us, PEB seems to be superior in green-light conditions.

Before the rise of photosynthesis and oxygen, Earth's oceans contained dissolved reduced iron (iron deposited in the absence of oxygen). Oxygen released by the rise of photosynthesis in the Archean eon then led to oxidized iron in seawater. The paper's computer simulations also found oxygen released by early photosynthesis led to a high enough concentration of oxidized iron particles to turn the surface water green.

Once all iron in the ocean was oxidized, free oxygen (0₂) existed in Earth's oceans and atmosphere. So a major implication of the study is that pale-green dot worlds viewed from space are good candidate planets to harbour early photosynthetic life.

Related: A mysterious 'black hole' in Pacific Ocean that sparked wild rumors online

The changes in ocean chemistry were gradual. The Archaean period lasted 1.5 billion years. This is more than half of Earth's history. By comparison, the entire history of the rise and evolution of complex life represents about an eighth of Earth's history.

Almost certainly, the color of the oceans changed gradually during this period and potentially oscillated. This could explain why blue-green algae evolved both forms of photosynthetic pigments. Chlorophyll is best for white light which is the type of sunlight we have today. Taking advantage of green and white light would have been an evolutionary advantage.

Could oceans change color again?

The lesson from the recent Japanese paper is that the color of our oceans are linked to water chemistry and the influence of life. We can imagine different ocean colors without borrowing too much from science fiction.

Purple oceans would be possible on Earth if the levels of sulphur were high. This could be linked to intense volcanic activity and low oxygen content in the atmosphere, which would lead to the dominance of purple sulphur bacteria.

Red oceans are also theoretically possible under intense tropical climates when red oxidized iron forms from the decay of rocks on the land and is carried to the oceans by rivers or winds. Or if a type of algae linked to "red tides" came to dominate the surface oceans.

These red algae are common in areas with intense concentration of fertilizer such as nitrogen. In the modern oceans, this tends to happen in coastline close to sewers.

As our sun ages, it will first become brighter leading to increased surface evaporation and intense UV light. This may favor purple sulphur bacteria living in deep waters without oxygen.

It will lead to more purple, brown, or green hues in coastal or stratified areas, with less deep blue color in water as phytoplankton decline. Eventually, oceans will evaporate completely as the sun expands to encompass the orbit of Earth.

At geological timescales nothing is permanent and changes in the color of our oceans are therefore inevitable.

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

We think of oxygen as life, sustenance, a literal breath of fresh air. But it's actually a very reactive element. Anyone who's burned a log has witnessed this firsthand. So why do so many life-forms breathe oxygen?

There are probably thousands of kinds of metabolisms, or chemical processes that maintain life, said Donald Canfield, a geobiologist at the University of Southern Denmark, but "virtually all eukaryotes" (life-forms whose cells contain a nucleus) and a vast array of prokaryotes (life-forms that lack a nucleus), use oxygen.

Canfield is talking primarily about heterotrophs — organisms, including humans, that get their nutrients and energy by consuming other organic matter. Not all organisms do this exclusively. For example, "plants get their carbon from CO2 in the air," said Dan Mills, a postdoctoral researcher at the University of Munich.

Heterotrophs break down organic matter in food by stripping electrons off of it. These are passed from one enzyme to another in the membrane of the mitochondria, generating a small current that pumps protons across this barrier. And given its high electronegativity, oxygen usually serves as the final station on this electron transport chain, accepting the electrons and picking up two protons to form water.

The process essentially creates a reservoir of protons that then flood through a protein channel in the membrane like a tiny hydroelectric dam. And, like a turbine, the protein synthesizes energy in the form of adenosine triphosphate (ATP) as it spins, explained Nick Lane, a professor of evolutionary biochemistry at University College London, in a public presentation. The cell can then use this packaged energy or send it off into the body to do things.

Life can use many other electron acceptors — like sulfate, nitrate and iron — but oxygen is the highest-energy acceptor available.

"The reduction of oxygen provides the largest free energy release per electron transfer, except for the reduction of fluorine and chlorine," University of Washington professor David Catling and his co-authors explained in a paper published in the journal Astrobiology.

Related: What is the world's most dangerous chemical?

Chlorine and oxygen can generate similar amounts of energy. Fluorine could certainly provide more energy than oxygen, but "fluorine is [...] useless as a biological oxidant because it generates an explosion upon contact with organic matter," they wrote in the study. That's not a gas you'd want to breathe.

Chlorine and fluorine are also poisonous, which highlights another benefit of oxygen. Aerobic respiration doesn't produce any toxic compounds, just water and carbon dioxide. However, oxygen's reactivity can be an issue if it builds up in tissues, where it can damage cellular components like DNA and proteins. That's why antioxidants, in moderation, are good for our health.

Oxygen is also far more abundant than fluorine, chlorine or the myriad electron acceptors used in other forms of respiration. Despite its proclivity for forming compounds with other atoms, a copious amount of oxygen is constantly produced via photosynthesis. This enables it to accumulate in the atmosphere and dissolve in water, where it is readily available to life. And, as a gas, it's easy to transport across membranes, Canfield and Mills explained.

Speaking of abundance, why not use nitrogen, which comprises 78% of Earth's atmosphere?

"The main problem with nitrogen is that it's triple bonded," Canfield said. "And it's very, very difficult to break."

Nitrogen is an important component of many biologic compounds, and there are whole groups of organisms that specialize in the energy-intensive processes required to break nitrogen's strong bonds to make it bioavailable, Canfield said.

Oxygen's unique utility comes down to quantum physics. Oxygen in its normal ground state can only accept electrons in the same spin state, not as an electron pair, which is the usual currency of chemistry.

"So the real trick to oxygen is that it can accumulate to high levels without reacting, but releases a lot of energy (to pump protons) when it is fed electrons one at a time," Lane told Live Science in an email.

So it seems oxygen sits in a sweet spot of reactivity and availability. It's milder than halogens such as chlorine and fluorine, and it isn't bound too strongly, like nitrogen. But it's much more reactive than other electron acceptors, like sulfate and nitrate.

Oxygen is easy to acquire, and it doesn't generate toxic compounds that require further processing. What's more, plants produce copious amounts of this reactive gas through photosynthesis, enabling us to use it to fuel our own bodies.

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

Scientists have developed a new method to break down plastic waste, using moisture from the air.

By exposing a common type of plastic to an inexpensive catalyst and leaving it exposed to ambient air, researchers broke down 94% of the material in just four hours.

The plastic transformed into terephthalic acid (TPA), a highly valuable building block for polyesters. Because TPA can be upcycled into more valuable materials, the process offers a safer and cheaper alternative to current plastic recycling methods. The researchers published their findings Feb. 3 in the journal Green Chemistry.

"The U.S. is the number one plastic polluter per capita, and we only recycle 5% of those plastics," co-corresponding author Yosi Kratish, a research assistant professor of chemistry at Northwestern University, said in a statement. "What's particularly exciting about our research is that we harnessed moisture from air to break down the plastics, achieving an exceptionally clean and selective process. By recovering the monomers, which are the basic building blocks of PET [polyethylene terephthalate], we can recycle or even upcycle them into more valuable materials."

Plastic waste is an increasingly important issue. Over half of the plastic ever made has been produced since 2000, and annual production is projected to double by 2050, according to the European Environment Agency.

To date, only 9% of the plastics ever produced have been recycled. The remainder, with lifetimes often lasting generations, can have serious environmental and health impacts. For example, they wash out to sea to form floating blobs of trash, harm wildlife, and break down into microplastics that can enter the human brain and other parts of our bodies.

Related: Will we ever be able to stop using plastic?

To find a new method to break down some of this waste, the researchers applied a molybdenum catalyst — a silver, ductile metal — and activated carbon to PET, the most common type of polyester plastic. The researchers then heated the mixture. After a short time, this broke the polyethylene's chemical bonds.

Then, when the team exposed the material to air, the mixture transformed into TPA, a valuable polyester precursor; and acetaldehyde, an industrial chemical that is also valuable and is easy to lift from the mixture.

When they tested the method on mixed plastics, the researchers found that it had an effect only on the polyester materials. That meant they didn't have to presort the plastics. It worked on plastic bottles, T-shirts and colored plastics, breaking them down into pure, colorless TPA.

"It worked perfectly," Kratish said. "When we added extra water, it stopped working because it was too much water. It's a fine balance. But it turns out the amount of water in air was just the right amount."

The team's next steps will be to adapt the process to large-scale industrial applications.

"Our technology has the potential to significantly reduce plastic pollution, lower the environmental footprint of plastics and contribute to a circular economy where materials are reused rather than discarded," study first author Naveen Malik, who was a researcher at Northwestern University at the time, said in the statement. "It's a tangible step toward a cleaner, greener future, and it demonstrates how innovative chemistry can address global challenges in a way that aligns with nature."

Scientists have made a cheap and flexible solar cell that lasts nearly 10 times longer than others of its type, an advance that could one day help to revolutionize solar energy production.

Often referred to as the "holy grail" of solar power, perovskite cells offer a lightweight alternative to traditional silicon-based solar technology. Their flexible structure enables them to be applied to cars and phones in the form of a printable layer so they can charge on the go.

Sounds too good to be true? So far, you're right. Perovskites come with some major flaws. Notably, they degrade quickly due to chemical reactions with moisture in the air that make them leak iodine.

But now, a team of researchers has found a solution to this problem. By embedding nanoparticles within the perovskites, they produced a new cell that lasts for 1,530 hours, a near-tenfold increase on previous perovskite solar cell designs. The researchers published their findings Feb. 20 in the journal EES Solar.

"By addressing these common challenges we see with perovskite solar technology, our research blows the doors wide open for cheaper, more efficient and more widely accessible solar power," study co-author Imalka Jayawardena, an engineering researcher at the University of Surrey's Advanced Technology Institute in the U.K., said in a statement. "What we've achieved here is a critical step toward developing high-performance solar cells that can withstand real-world conditions — bringing us closer to their commercial use at a global scale."

Solar power surge

As the fastest-growing and cheapest form of renewable energy, solar power is key to cutting greenhouse gas emissions. But the technology's growth is hampered by its reliance on silicon, a finite and non-renewable resource that, in its purest form, is costly to produce.

To get around this bottleneck, scientists have looked to develop perovskite alternatives — synthetic versions of naturally occurring calcium titanium oxide crystals that can be made at a fraction of the cost. But unlike pure silicon cells, which can last for decades, solar cells made from perovskite only last for 100 or so hours, drastically limiting their utility.

Related: Ultra-thin solar 'coating' can turn phone cases and EVs into mini power generators

In the new study, the scientists looked for a way to trap the iodine that leaks from perovskites. Their solution was to embed tiny nanoparticles of aluminum oxide within the cells as they were manufactured. This not only prevented the iodine from leaking but also created a more uniform and electrically conductive structure.

After testing these cells under extreme heat and humidity, the researchers found that the modified cells maintained a high performance for more than two months (1,530 hours), a significant improvement on the 160-hour lifespan of unenhanced perovskite cells.

The researchers plan to continue investigating their new technique to see if these gains can be built upon further.

"A decade ago, the idea of perovskite solar cells lasting this long under real-world conditions seemed out of reach," study lead author Hashini Perera, a researcher at the Advanced Technology Institute, said in the statement. "With these improvements, we're breaking new ground in stability and performance, bringing perovskite technology closer to becoming a mainstream energy solution."

For the first time, scientists have observed water molecules splitting in real time to form hydrogen and oxygen.

And right before they split, the molecules did something completely unexpected: They flipped 180 degrees.

This micro acrobatic stunt takes energy, which offers a crucial explanation for why splitting water takes more energy than theoretical calculations suggested.

The researchers say that studying this further could offer key insights into making the process of splitting water molecules more efficient — opening a pathway to cheaper clean hydrogen fuel and breathable oxygen for future Mars missions. They published their findings March 5 in the journal Science Advances.

Making hydrogen fuel

Hydrogen has a number of key properties that make it an enticing source of green energy. The energy-rich fuel is capable of powering trucks and even cargo ships, and it is the only alternative to fossil fuels in industries such as steel and fertilizer manufacturing. When it's burned, the fuel releases water instead of carbon dioxide.

Yet the steep energy requirements for hydrogen production severely limit the scale at which the fuel is produced. According to the International Energy Authority, 322 million tonnes (354 million tons) of hydrogen fuel needs to be produced each year to meet global energy needs. But in 2023, only 97 million tonnes (107 million tons) was manufactured at a monetary cost 1.5 to six times greater than fossil fuel production — and the vast majority of it was made using fossil fuels too.

Related: Scientists discover revolutionary method that makes fuel from water and sunlight — but it's not finished yet

Hydrogen fuel is made by adding water to an electrode and then splitting the water with an applied voltage into hydrogen and oxygen.

This process is most efficient when the chemical element iridium is used as a catalyst for the oxygen evolution reaction that cleaves oxygen from water molecules. But iridium only arrives on our planet from meteorite impacts, making it costly and scarce.

But even when using iridium, the process is less efficient than scientists believe it should be.

"It ends up taking more energy than theoretically calculated. If you do the math, it should require 1.23 volts. But, in reality, it requires more like 1.5 or 1.6 volts," study lead author Franz Geiger, a professor of chemistry at Northwestern University, said in a statement. "Providing that extra voltage costs money, and that's why water splitting hasn't been implemented at a large scale."

To better understand the energy requirements of this process and why it's less efficient than theory suggests, the researchers placed water on an electrode inside a container and measured the molecules' positions using the amplitude and phase of laser light shone onto them.

When the scientists applied a voltage across the electrode, they observed that the molecules rapidly flipped and rotated so that their two hydrogen atoms touching the electrode faced up and the oxygen atom faced down.

"Electrodes are negatively charged, so the water molecule wants to put its positively charged hydrogen atoms toward the electrode's surface," Geiger said. "In that position, electron transfer from water's oxygen atom to the electrode's active site is blocked. When the electric field becomes strong enough, it causes the molecules to flip, so the oxygen atoms point toward the electrode's surface. Then, the hydrogen atoms are out of the way, and the electrons can move from water's oxygen to the electrode."

By measuring the number of molecules that rotated and the energy required for them to do so, the researchers found that this flipping was likely a necessary and unavoidable part of the splitting process. What's more, the researchers discovered that higher pH levels made this process more efficient.

Further studying this process could help scientists to design more efficient catalysts to use in the process, and to better understand the chemical processes involved, the researchers said, while also offering fresh insights into how water behaves.

"Our work underscores how little we know about water at interfaces," Geiger said. "Water is tricky, and our new technology could help us understand it a bit better."

"By designing new catalysts that make water flipping easier, we could make water splitting more practical and cost-effective," he added.

Scientists have invented a smart fabric that converts light into heat and can raise temperatures by more than 54 degrees Fahrenheit (30 degrees Celsius) after just 10 minutes in the sun. The new material could be used in clothing designed for very cold temperatures.

Specialized nanoparticles that absorb sunlight and convert it to heat are embedded within the new material, which was described late last year in the journal Advanced Composites and Hybrid Materials. At the same time, temperature-responsive dyes incorporated into the fibers reversibly change color, allowing users to visually monitor temperature fluctuations.

Maintaining body temperature

For years, scientists have designed wearable heaters to help maintain a comfortable body temperature in cold environments. Such fabrics could be used in mountain rescue equipment and even pet clothing, but existing designs typically rely on expensive components such as metal nanomaterials or cumbersome battery-powered heating elements.

To get around these problems, chemical engineer Yuning Li and his team at the University of Waterloo in Canada looked to photothermal polymers, which are plastic-like materials that convert light into heat.

Nanoparticles of the two polymers — polyaniline (PANI) and polydopamine (PDA) — are embedded within a matrix of thermoplastic polyurethane (PTU) fibers, a material widely used to produce waterproof clothing and sportswear. The team also incorporated various temperature-responsive (thermochromic) dyes into the mix during the spinning process, producing a series of fibers that changed color as the temperature of the material increased.

These newly spun fibers were readily woven into fabric and the team knitted a tiny sweater for a teddy bear to test the properties of the smart material. The red jumper reached an impressive 128.3 F (53.5 C) after just 10 minutes of sun exposure. As the temperature climbed, the red dye molecules changed chemical structure, causing them to turn white.

"The incorporated nanoparticles are highly efficient at absorbing sunlight across a range of wavelengths," Li told Live Science in an email. "When sunlight hits these nanoparticles, they absorb the energy and release it as heat through a process called photothermal conversion."

The smart fabric has a soft and elastic texture, which allows the material to stretch by as much as five times its original size and retain its color- and temperature-changing properties even after 25 washes, according to the study. "We prioritized durability, ensuring the fabric could withstand repeated use and environmental exposure while maintaining its innovative properties," Li said.

The team is working to prepare the material for commercial manufacturing, but they still have to do further testing before it can gain widespread use.

"The next steps for this research focus on reducing production costs, scaling up the fabrication process, and ensuring the fibers are safe for prolonged skin contact," Li said.

Editor's Note: The headline for this story was updated on Monday, Feb. 3 at 1:11 p.m. EST to note that the fabric heats up more than 50 degrees, not almost 50 degrees.

Scientists have used artificial intelligence (AI) to design never-before-seen nanomaterials with the strength of carbon steel and the lightness of styrofoam.

The new nanomaterials, made using machine learning and a 3D printer, more than doubled the strength of existing designs. The scientists behind the new study said they could be used in stronger, lighter and more fuel-efficient components for airplanes and cars. They published their findings Jan. 23 in the journal Advanced Materials.

"We hope that these new material designs will eventually lead to ultra-light weight components in aerospace applications, such as planes, helicopters and spacecraft that can reduce fuel demands during flight while maintaining safety and performance," co-author Tobin Filleter, a professor of engineering at the University of Toronto, said in a statement. "This can ultimately help reduce the high carbon footprint of flying."

In many materials, strength and toughness can often be at odds. Take a ceramic dinner plate, for example: while plates are usually strong and can carry heavy loads, their strength comes at the cost of toughness — it doesn’t take much energy to make them shatter.

Related: Scientists discover revolutionary method that makes fuel from water and sunlight — but it's not finished yet

The same problem applies to nano-architectured materials, whose construction from multitudes of tiny, repeating building blocks 1/100th the thickness of a human hair makes them strong and stiff for their weight, but can also cause stress concentrations that lead to sudden breakages. So far, this tendency to shatter has limited the materials' applications.

"As I thought about this challenge, I realized that it is a perfect problem for machine learning to tackle," first-author Peter Serles, an engineering researcher at Caltech, said in the statement.

To search for better ways to design nanomaterials, the researchers simulated possible geometries for their design before passing them through a machine learning algorithm. By learning from the designs they had generated, the algorithm was able to predict the best shapes that would evenly distribute applied stresses while also carrying a heavy load.

With these shapes in hand, the researchers used a 3D printer to create their new nanolattices, finding that they could withstand a stress of 2.03 megapascals for every cubic meter per kilogram — a strength five times higher than titanium.

"This is the first time machine learning has been applied to optimize nano-architected materials, and we were shocked by the improvements," Serles said. "It didn’t just replicate successful geometries from the training data; it learned from what changes to the shapes worked and what didn’t, enabling it to predict entirely new lattice geometries."

The researchers said their next steps will center on scaling up the materials until they can be used to make bigger components, while also searching for even better designs using their process. The primary aim is to design much lighter and stronger components for vehicles in the future.

"For example, if you were to replace components made of titanium on a plane with this material, you would be looking at fuel savings of 80 litres per year for every kilogram of material you replace," Serles said.

Scientists have discovered the essential building blocks to life on a sample from a distant asteroid.

The sample, which the OSIRIS-REx spacecraft collected from the asteroid Bennu and returned to Earth in 2023, contains all five nucleobases — the "letters" that make up DNA and RNA — alongside mineral compounds, all of which have never previously been seen on extraterrestrial rocks.

The minerals are rich in carbon, sulfur, phosphorus, fluorine and sodium, making them resemble those left in the crusts of dried lake beds on Earth — except they date to the birth of the solar system 4.6 billion years ago. These elements, alongside the five nucleobases that make up DNA and RNA, are the basic building blocks for life on our planet.

The two teams of researchers who made the discoveries published their findings Jan. 29 in two papers in the journal Nature Astronomy.

"We now know from Bennu that the raw ingredients of life were combining in really interesting and complex ways on Bennu's parent body," study co-lead author Tim McCoy, curator of meteorites at the Smithsonian's National Museum of Natural History, said in a statement. "We have discovered that next step on a pathway to life."

Bennu is a potentially hazardous asteroid that has a 1-in-2,700 chance of striking Earth in the year 2182 — the highest odds of any known space object. But scientists are more interested in what's trapped on the space rock: As a carbon-rich asteroid, it likely contains many of the primordial molecules present when life first emerged on Earth.

Related: NASA's OSIRIS-REx mission almost bit the dust — then Queen guitarist Brian May stepped in

OSIRIS-REx launched in September 2016 and traveled 200 million miles (320 million kilometers) to reach Bennu.

Once there, the spacecraft orbited the asteroid for nearly two years as flight engineers searched for a landing site. Upon making contact with the space rock, OSIRIS-REx fired a burst of nitrogen from its Touch-and-Go Sample Acquisition Mechanism to stick the landing and prevent itself from sinking through the asteroid. The nitrogen burst captured a 4.29-ounce (121.6 grams) sample in the process.

In October 2023, the sample was brought to Earth aboard OSIRIS-REx's capsule, which reached speeds of up to 27,000 mph (43,000 km/h) before it deployed its parachute and landed safely in the Utah desert. To avoid contamination, the sample container was taken to a clean room before being opened.

The researchers behind the first study received slices of the Bennu sample, which they examined under a scanning electron microscope. This enabled the team to study features on the sample's surface with a resolution of one-hundredth the width of a human hair.

The scientists discovered sodium carbonate, typically found in evaporated lakes that once contained life on Earth, on the space rock's surface. Within the sodium carbonate, the team found 11 minerals that are important precursors for organic compounds. The mineral compositions differed subtly from those found on our planet; being rich in phosphorus and poor in boron, when the reverse is true in Earth’s lakes.

The researchers think brine similar to that found on Bennu could also exist on other bodies in the solar system, such as the dwarf planet Ceres and Saturn's icy moon Enceladus.

In the second study, conducted by scientists in Japan, a separate piece of the sample was also found to contain the five nucleobases — adenine, guanine, cytosine, thymine and uracil — which combine with ribose and phosphate to form DNA and RNA, the ladder-like structures that make up the genetic code of all life on Earth.

This is the first time that scientists have found these nucleobases on a distant asteroid. In 2023, a sample taken from the Ryugu space rock by the Hayabusa2 spacecraft was found to contain uracil, yet the other nucleobases were missing.

It's unclear what this means for life beyond our planet. While the existence of these minerals on Bennu is a sure indication that the asteroid had the right ingredients for life, the researchers are unsure if the asteroid's environment was too harsh for the compounds to grow into complex organic structures. To further investigate, the scientists plan to reexamine meteorites in their collection for similar salts and compounds.

"We now know we have the basic building blocks to move along this pathway towards life, but we don't know how far along that pathway this environment could allow things to progress," McCoy said.

Chemists have invented a new material that could be the future of body armor — chainmail. But this isn't the Middle Ages all over again; the new super-strong material is made of molecules that are interlocked on a nanoscale, scientists say.

Researchers fused lines of molecules like links in a chain to create sheets of the world's first two-dimensional mechanically interlocked material (2D MIM), which has length and width. The material contains 100 trillion chemical bonds per square centimeter (around 650 trillion per square inch), which is the highest density of mechanical bonds ever achieved, the researchers reported in the study, published Jan. 16 in the journal Science.

The study authors added a small amount of the material to a tough plastic material called Ultem — also made from molecule chains. Ultem is already incredibly strong but became even stronger with the 2D MIM. The research, which could eventually be used in body armor, was partly funded by the government's Defense Advanced Research Projects Agency.

"It's similar to chainmail in that it cannot easily rip because each of the mechanical bonds has a bit of freedom to slide around," study co-author William Dichtel, a chemistry professor at Northwestern University in Illinois, said in a statement. "If you pull it, it can dissipate the applied force in multiple directions. And if you want to rip it apart, you would have to break it in many, many different places."

Related: Scientists discover revolutionary method that makes fuel from water and sunlight — but it's not finished yet

The 2D MIM material is made of interlocked polymers, which are long chains of smaller molecules, called monomers. The team took lines of X-shaped monomers and arranged them into crystal structures that react together so that the ends of the monomers bond with the ends of other monomers, according to the statement.

Each monomer's X-shape left gaps in which researchers could weave additional lines of these molecular building blocks, creating layers of interlocked 2D polymers within the crystals. The scientists then dissolved the crystals to retrieve the interlocked polymers.

"After the polymer is formed, there's not a whole lot holding the structure together," Dichtel said. "So, when we put it in solvent, the crystal dissolves, but each 2D layer holds together. We can manipulate those individual sheets."

To test their new material, the researchers made composite materials out of 97.5% Ultem fiber and 2.5% 2D MIM. The small amount of interlocked 2D polymer increased the force needed to deform Ultem fibers by 45% and the amount of stress the Ultem could withstand by 22%, according to the study.

"We have a lot more analysis to do, but we can tell that it improves the strength of these composite materials," Dichtel said. "Almost every property we have measured has been exceptional in some way."

Scientists in Japan have demonstrated a new proof-of-concept reactor that can harvest renewable hydrogen fuel from sunlight and water.

The new 1,076-square-foot (100 square meters) reactor uses photocatalytic sheets to split apart the oxygen and hydrogen atoms found in water molecules, thus siphoning the hydrogen away to be used as fuel.

While the technology remains in its infancy, the scientists behind the research say that, if more efficient photocatalysts can be developed, their breakthrough could enable the production of cheap, sustainable hydrogen fuel to meet various energy needs. They published their findings Dec. 2 in the journal Frontiers in Science.

"Sunlight-driven water splitting using photocatalysts is an ideal technology for solar-to-chemical energy conversion and storage, and recent developments in photocatalytic materials and systems raise hopes for its realization," senior author Kazunari Domen, a chemistry professor at Shinshu University in Japan, said in a statement. "However, many challenges remain."

Upon being exposed to light, photocatalysts boost chemical reactions that break water molecules down into their constituent parts. However, most existing "one-step" catalysts — which decompose water into hydrogen and oxygen in one go — are extremely inefficient, leaving most of the hydrogen fuel to be refined using natural gas, a fossil fuel.

Related: 'Holy grail' of solar technology set to consign 'unsustainable silicon' to history

To look for a way past this deadlock, the researchers behind the new study investigated a photocatalyst that uses a more sophisticated two-step process, with one step separating out the oxygen and the next step removing the hydrogen.

Creating a photocatalyst for this process enabled the scientists to build their prototype reactor, which ran for three years and worked even better using real sunlight than the ultraviolet light used in the lab.

"In our system, using an ultraviolet-responsive photocatalyst, the solar energy conversion efficiency was about one and a half times higher under natural sunlight," first author Takashi Hisatomi, a researcher at Shinshu University, said in the statement. "Simulated standard sunlight uses a spectrum from a slightly high latitude region. In an area where natural sunlight has more short-wavelength components than simulated reference sunlight, the solar energy conversion efficiency could be higher."

Despite these promising gains, the efficiency of the reaction is still too low for commercial use.

"Currently, the efficiency under simulated standard sunlight is 1% at best, and it will not reach 5% efficiency under natural sunlight," Hisatomi said.

To make the important strides to increase efficiency, the scientists have called on others to create better photocatalysts and larger reactors. Work on safety will also be vital: Hydrogen fuel refining also produces the explosive byproduct oxyhydrogen, which can be safely disposed of in the two-step process.

"The most important aspect to develop is the efficiency of solar-to-chemical energy conversion by photocatalysts," Domen said. "If it is improved to a practical level, many researchers will work seriously on the development of mass production technology and gas separation processes, as well as large-scale plant construction. This will also change the way many people, including policymakers, think about solar energy conversion, and accelerate the development of infrastructure, laws, and regulations related to solar fuels."