In a development that could reshape how chemists recycle one of the most persistent classes of synthetic molecules ever made, researchers have reported a new catalytic strategy that strips fluorine atoms from perfluoroalkyl chains in a controlled, sequential fashion while installing a single boron group at the chain’s terminus. The work, published in Nature Communications, introduces what the authors call zipper polydefluorination-monoborylation, a transformation that walks down a perfluoroalkyl skeleton and removes multiple fluorine atoms in a stepwise cascade before capping the partially defluorinated product with a boron substituent. The achievement addresses a longstanding challenge in fluorine chemistry: while carbon-fluorine bonds are prized in materials and pharmaceuticals for their extraordinary stability, that same stability makes them notoriously difficult to break selectively, leaving chemists with few tools for converting perfluorinated waste streams into useful building blocks.

Perfluoroalkyl substances, often grouped under the umbrella of PFAS chemistry, owe their remarkable properties to the strength of the carbon-fluorine bond, which at roughly 485 kilojoules per mole in fully fluorinated chains ranks among the strongest single bonds that carbon forms. This bond strength, combined with the sheer electronegativity of fluorine, means that perfluoroalkyl groups resist oxidation, hydrolysis, and thermal degradation under nearly all conventional conditions. As a consequence, perfluorinated chains accumulate in the environment and resist standard degradation pathways, prompting intensive research into defluorination methods ranging from photocatalytic single-electron transfer to reductive hydrodefluorination with metal hydrides. Yet most existing approaches suffer from a common limitation: they either remove fluorine indiscriminately, collapsing the entire chain to a mixture of products, or they remove only a single fluorine atom, offering little control over the degree of defluorination. The new study demonstrates a way to achieve the best of both worlds, delivering multi-fluorine removal in a predictable pattern while simultaneously introducing a versatile functional handle.

The conceptual heart of the method lies in what the authors describe as a zipper mechanism. Rather than removing fluorine atoms at random positions along the chain, the reaction initiates at a defined site and then propagates along the perfluoroalkyl backbone one carbon at a time, much as a zipper opens tooth by tooth. At each step along the chain, a fluorine atom is expelled and replaced in sequence, with the reaction front advancing until a thermodynamic or kinetic stopping point is reached. The monoborylation component of the process ensures that once the zipper has run its course, the resulting olefinic or anionic intermediate is intercepted by a boron reagent, furnishing a product in which the boron group occupies a predictable position derived from the original perfluoroalkyl terminus. Organoboron compounds are among the most valuable intermediates in modern synthetic chemistry, serving as partners in Suzuki-Miyaura cross-coupling and as precursors to alcohols, amines, and halides, so their installation at a defluorination site effectively converts a chemical dead end into a synthetic launching pad.

Mechanistically, the transformation draws on the characteristic reactivity of organoboron reagents toward geminal difluoroalkenes, the partially defluorinated intermediates that arise when consecutive fluorine atoms are eliminated from a perfluoroalkyl chain. Perfluoroalkyl chains, when subjected to strong bases or reductive conditions, can undergo sequential beta-fluoride elimination events, each generating a more highly unsaturated species. The boron reagent in the reported system engages these unsaturated intermediates through addition pathways that install boron at the terminal carbon while generating a carbanionic center adjacent to the remaining difluoromethylene unit. That newly formed carbanion is perfectly positioned to expel another fluoride ion, regenerating an unsaturated handle one carbon further down the chain and setting up the next iteration of the cycle. This intimate coupling of borylation and defluorination is what gives the reaction its chain-walking character: each borylation event creates the conditions for the next defluorination event, and the process ripples along the chain until the fluorine reservoir is exhausted or the reagents are consumed.

The practical significance of the work becomes apparent when one considers the fate of the fluorine atoms themselves. In the zipper process, fluorine is released as fluoride ion, which under the reaction conditions is sequestered by silicon-containing trapping agents to form robust silicon-fluorine bonds. Fluorosilicates of this kind are stable, isolable, and, importantly, compatible with industrial processes for recycling fluorine back into hydrofluoric acid, the feedstock from which essentially all synthetic organofluorine compounds are made. In this sense, the method does more than transform perfluoroalkyl chains into borylated products; it also lays the groundwork for closing the fluorine cycle, recovering the fluorine content of persistent molecules in a form that can be fed back into the chemical economy. With global concern over PFAS contamination intensifying and regulatory pressure on fluorinated chemicals mounting, technologies that enable fluorine recovery rather than dilution into the environment carry considerable weight.

From a synthetic chemistry standpoint, the products of the reaction open multiple downstream avenues. Because the boron substituent is installed at a defined position on a partially fluorinated alkyl or alkenyl framework, the resulting molecules retain a programmable mixture of carbon-fluorine bonds and carbon-boron functionality. The remaining fluorine atoms can tune the lipophilicity, metabolic stability, and dipole moment of the scaffold, properties that are actively exploited in medicinal chemistry when designing fluorinated drug candidates. Meanwhile, the boron group serves as a convertible handle, allowing the defluorinated carbon center to be joined to aromatic rings through cross-coupling, oxidized to alcohols, or aminated in homologation reactions. In effect, the method converts a chemically inert perfluoroalkyl group, useful mainly as a passive coating or tail, into a functionalized partial-fluoro skeleton that can participate in the full toolkit of carbon-carbon and carbon-heteroatom bond formation.

The reported study also sheds light on the factors governing how far the zipper travels along a given chain. Chain length, the nature of the substituent anchoring the perfluoroalkyl group, and the stoichiometry of the base and boron reagent all influence the number of fluorine atoms removed and the distribution of products obtained. Through systematic variation of these parameters and careful analysis of the reaction mixtures by nuclear magnetic resonance spectroscopy and mass spectrometry, the authors map out the scope of the transformation across a range of perfluoroalkyl substrates, from short perfluoroethyl groups to longer chains approaching those found in commercial fluorosurfactants. Control experiments interrogating individual elementary steps support the proposed sequential elimination-addition mechanism and rule out alternative pathways involving radical fragmentation or random beta-scission, underscoring the unusually high degree of positional control the method achieves.

Beyond its immediate synthetic utility, the work contributes to a growing body of research aimed at taming the carbon-fluorine bond through the strategy of defluorinative functionalization. Over the past decade, chemists have developed an impressive repertoire of reactions that replace fluorine in perfluoroalkyl and fluoroarene substrates with groups ranging from hydrogen to sulfur, nitrogen, and carbon substituents. What distinguishes the zipper approach is its multiplicity and its directionality: instead of a single substitution event at a single site, the reaction performs many substitutions in a choreographed sequence along a chain, effectively treating the perfluoroalkyl group as a one-dimensional lattice to be patterned at will. This conceptual reframing, from defluorination as bond-breaking to defluorination as programmable editing, may inspire analogous strategies for other polyfluorinated motifs, including perfluoroarenes and perfluoropolyethers, two additional classes of highly persistent fluorinated materials.

Challenges remain before such chemistry can move from the laboratory bench to environmental remediation at scale. The reaction conditions reported rely on strong bases and moisture-sensitive boron reagents handled under inert atmosphere, and the substrates studied are discrete, soluble perfluoroalkyl compounds rather than the polymeric or adsorbed PFAS species that dominate environmental contamination. Scaling the chemistry will require sourcing boron reagents economically, managing the large quantities of fluoride generated, and integrating the transformation with upstream methods for extracting perfluoroalkyl pollutants from water and soil. Nevertheless, the demonstration that a perfluoroalkyl chain can be unzipped in a controlled manner, with fluorine captured in recyclable form and a valuable boron handle installed, establishes a credible chemical pathway toward defluorination that is both productive and selective, a combination that has eluded the field for decades.

As policymakers and scientists grapple with the legacy of fluorine chemistry, work of this kind illustrates a middle path between the wholesale abandonment of fluorinated molecules and their indefinite persistence in the biosphere. The fluorine atom confers properties that industry and medicine continue to depend upon, and the ability to disassemble fluorinated architectures on demand, recovering both the carbon skeleton as a functionalized building block and the fluorine content as a recyclable salt, points toward a more circular future for organofluorine chemistry. The zipper polydefluorination-monoborylation strategy reported in Nature Communications is an early but striking example of that vision made concrete at the molecular level, and it is likely to prompt rapid efforts to extend the concept across the breadth of persistent fluorinated chemicals.

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Author: José Luis Granados Mateo is a postdoctoral researcher in the Department of Philosophy at the University of the Basque Country (EHU) and a member of the Integrated History and Philosophy of Science (iHPS) research group. His work focuses on history and philosophy of science, science and values, and the epistemology of scientific practices.

By 1800, water had already changed once. It was no longer one of nature’s simple terms, but a compound of hydrogen and oxygen (see Is Water H₂O? (1) From Element to Compound). The new chemistry had a powerful account of that change. What it still wanted was a more direct experiment: water itself, giving up its constituents.

Then came Volta’s pile.

It was not a spectacular object. Discs of metal stacked in pairs, pieces of moist material between them. Yet it altered the experimental life of electricity. Static machines produced sparks, shocks and theatrical moments. The pile offered something steadier. Electricity could now be made to act for long enough to enter chemistry.

Nicholson and Carlisle tried it on water almost at once. They placed wires from the pile in the liquid. Gas appeared, or rather the two sides of the reaction began to declare themselves. Hydrogen was released at one wire; oxygen, in the earliest arrangement, fixed itself to the other, nearly two inches away. Later, with gold or platinum electrodes, oxygen could be collected as gas. The awkward fact remained. The two supposed products of one decomposition appeared at opposite electrical ends of the apparatus.

At first glance, this looked like the experiment the new chemistry had been waiting for.

The trouble was that it looked a little too good.

A cleaner decomposition

The first article in this series followed water from element to compound. That change had not been a simple passage from error to fact. It involved a wider transformation of chemical practice: new gases, new names, balances, combustion experiments, recomposition, and a growing confidence in quantitative composition.

Electrolysis seemed to bypass much of that history. No flame, no combustion vessel, no burning of inflammable air with oxygen, no droplets collected on glass. The battery appeared to make water surrender its constituents directly. If water was composed of hydrogen and oxygen, here was synthesis in reverse: pass electrical action through the liquid and the two gases would come out.

Strictly speaking, electrical decomposition of water was not new. In 1789, Adriaan Paets van Troostwijk and Jan Rudolph Deiman had used repeated sparks of static electricity to obtain a mixture of hydrogen and oxygen from water, then recombined that mixture into water. Their experiment mattered. It was also awkward. The gases came together, in small quantities, and separating them for testing was difficult.

Nicholson and Carlisle’s version looked sharper. The Voltaic pile produced sustained action. The products could be made to appear separately. Hydrogen here, oxygen there. Each could be collected, examined and identified.

That was the improvement. It was also where the puzzle began.

When neatness becomes suspicious

Suppose water is decomposed. What, exactly, has been decomposed? A particle of water, perhaps. Yet if that particle is split into hydrogen and oxygen, why do the products not appear together? Why does hydrogen emerge at one wire and oxygen at the other? Why should the layout of the battery govern the place where each product appears, while the water between the wires shows no visible sign of being torn apart?

This was not an experiment going wrong. The gases appeared reliably. They appeared at the expected poles. That was precisely what made the situation uncomfortable.

Nicholson noticed the difficulty at once. Hydrogen appeared at one wire; oxygen was found at the other, nearly two inches away. The space between them did not look like a region where water was being visibly pulled into its parts. No trail of bubbles marked the route. The products simply appeared at opposite ends of the electrical arrangement, as if the battery had assigned them places.

Others saw the same embarrassment. Cuvier put it bluntly: if oxygen and hydrogen came from the same particle of water, why did they appear at two distant points, each always at its own wire? Haüy pressed the problem in molecular terms. If one molecule of water was decomposed, why were its products separated? If two molecules were decomposed, why did one give only hydrogen and the other only oxygen?

The question was almost childishly simple. That was why it would not go away.

Today we speak of ions, electrodes and charge transport. Those words make the old difficulty vanish before we have properly felt it. Around 1800, chemists had poles, attractions, affinities, particles, electrical fluids, acids, alkalis, metals, and a growing suspicion that electricity and chemistry belonged together. The experiment was clear enough. The mechanism was not.

Chang calls this the distance problem. The phrase is exact. The gases appeared with too much order for a simple story of local decomposition.

The distance problem

The obvious interpretation could still be defended. Perhaps electricity pulled the constituents of water apart. Oxygen went to one pole, hydrogen to the other. Many experiments already suggested that substances had different electrical tendencies, so the idea was hardly absurd.

It left the hardest part untouched.

What travelled through the water? Hydrogen? Oxygen? Something already separated? Something charged? If the constituents were detached somewhere between the electrodes, why did they not appear on the way? If they moved invisibly, why did they not recombine? And if decomposition happened only at the wires, how could it be the decomposition of the same portion of water?

Changes in apparatus did not remove the difficulty. Davy tried water held in separate gold cups, connected only by a moist bridge of fine asbestos, an arrangement he attributed to William Hyde Wollaston. The products still appeared at distant sites. The more carefully experimenters separated the places of action, the more pressing the question became: where, exactly, was water being decomposed?

That was the oddity. Electrolysis was not messy. If anything, it was too orderly. Hydrogen went one way, oxygen another. The battery gave a geography to the products of decomposition.

A flame did nothing like this. When inflammable air burned with oxygen, water appeared as the product of combination. There were difficulties, of course, but not difficulties of distance. Electrolysis changed the spatial form of chemical evidence. Water did not answer in one place. It answered at two.

Under electrical action, water no longer seemed merely something to be composed or decomposed. It seemed to contain routes.

Ritter changes the verb

Johann Wilhelm Ritter changed the verb.

For Lavoisierian chemists, electrolysis was analysis. Electricity decomposed water into hydrogen and oxygen. Ritter proposed the reverse. Electrolysis was synthesis.

At the negative pole, water combined with negative electricity and yielded hydrogen. At the positive pole, water combined with positive electricity and yielded oxygen. No product needed to travel from a decomposed particle somewhere in the middle of the liquid. The gases were made where they appeared.

It was a radical move. It was not foolish. Ritter took the most striking feature of the experiment seriously: the gases emerged at different poles. He did not treat that separation as an inconvenience to be patched up later. He built his interpretation around it. If hydrogen appears here and oxygen there, perhaps the two gases are being formed locally by two different electrical processes.

The price was high. Water became elementary again. Hydrogen and oxygen became compounds of water and electricity.

To modern chemistry, this looks perverse. Seen from the apparatus, it had one plain advantage: the gases were produced where they were found. Electricity was still widely treated as something with material agency. If a pile drove something into a liquid, why could that something not combine chemically with the liquid? Why must electricity be only the knife, and never part of what is made?

The phlogistonist echo was clear enough. If inflammable air could be read as water combined with phlogiston, hydrogen could be read as water combined with negative electricity. Identify phlogiston with negative electricity, and Ritter’s scheme comes surprisingly close to older anti-Lavoisierian accounts of water.

Priestley, from America, also saw that the new experiment did not force the Lavoisierian reading. He followed the debate through Nicholson’s journal and argued that electrolysis need not be understood as the decomposition of compound water. His intervention matters because it shows how differently the same electrical process could be read. For Lavoisierians, the gases were constituents released from water. For Priestley, they could still be products of another chemical-electrical process.

Ritter’s view did not last. Chemistry was becoming more firmly Lavoisierian, and elementary water had little room left in the emerging discipline. Still, electrolysis itself had not supplied a neat refutation. Ritter had solved the distance problem rather well. His view lost ground as chemists built an electrochemical practice around compound water.

Keeping water a compound

The Lavoisierian problem was delicate. Chemists had to hold three claims together. Water was composed of hydrogen and oxygen. Electricity decomposed it. Yet the products appeared at different electrodes, with no visible decomposition in between.

So the theory needed help.

One option was invisible transfer: one constituent moved through the liquid to its proper pole. Another was a chain of decompositions and recompositions, in which neighbouring particles handed hydrogen and oxygen along until the products emerged at the ends. In related versions, the whole liquid was imagined as electrically ordered, so that the final gases appeared only where that order met the metal.

These were rescue hypotheses. The phrase can sound dismissive, although it need not be. They were attempts to make a powerful practice intelligible. Electrolysis worked. Its effects could be repeated, varied and extended. What was missing was a mechanism that made the spatial separation of the gases sit comfortably with compound water.

The chain idea had a certain elegance. No single particle of hydrogen or oxygen needed to travel the whole distance. A constituent could be passed along through a succession of exchanges, like a message moving through a crowd. The gas would appear only at the end of the chain.

But what exactly was being passed? How were the exchanges coordinated? Why did the process end at just the right pole? The explanation helped, though it left plenty in the dark.

The transfer idea had its own attraction. If chemical particles had electrical characters, perhaps they could be drawn through the liquid. Again, however, the path remained hidden. The water between the wires did not display the process one wanted to see. Without a developed theory of ions and charge transport, the mechanism remained speculative.

None of this made the hypotheses useless. They kept research going. They suggested new apparatuses, new distances, new electrodes, new liquids and new comparisons. An experiment need not be fully understood before it becomes useful.

That is more or less what happened.

When water stopped being negotiable

By the early nineteenth century, compound water was becoming part of the working furniture of chemistry. Not every detail had been explained. The distance problem remained. Yet fewer and fewer chemists were prepared to reopen the older question of whether water might be elementary.

Water as a compound was no longer simply a hypothesis under trial. Increasingly, it was a condition for further work.

That shift mattered. Once water was taken as compound, electrochemistry could ask other questions. What moved towards each pole? What kinds of substances were decomposed by the pile? How did acids, alkalis, salts and metals behave under electrical action? What changed when the electrode, distance or solution changed? How much chemical change corresponded to how much electrical action?

The material practice hardened. Trough batteries became easier to handle. Gold and platinum electrodes helped prevent the products from disappearing into reactions with the metal. Separate cups connected by wet asbestos became a standard arrangement. The theory remained unsettled, but the apparatus was becoming disciplined.

The unresolved puzzle did not stop electrochemistry. Experimenters varied the arrangement. They separated vessels, introduced bridges, changed electrodes, added acids and salts, and studied what appeared where. The field learnt to handle the phenomena before it possessed a settled account of what moved inside the liquid.

Later developments brought more discipline. Faraday gave electrolysis a quantitative discipline in the 1830s. By the end of the nineteenth century, Arrhenius’s theory of ionic dissociation would help make sense of processes that had long remained obscure. That later clarity, though, should not be read backwards into 1800. For decades, electrochemistry advanced with a problem at its centre.

What electricity did to water

Electrolysis changed what water could be made to do in the laboratory. It did not simply reveal a truth waiting inside the liquid. It altered the route by which that truth could be pursued.

The first article in this series followed water into the balance sheet of chemistry. Water became a substance that could be produced, analysed, recomposed, weighed and named through its constituents. Electrolysis added another sort of depth. Water could now be interrogated electrically. Its constitution became a problem of poles, currents, attractions, transport and hidden motion.

Lavoisierian chemistry had given water constituents. Electrochemistry gave it pathways.

The familiar classroom demonstration hides that difficulty. The bubbles are easy to see. Their meaning was not. Hydrogen here, oxygen there: that simple separation forced chemists to imagine invisible processes inside a liquid that otherwise looked continuous and unchanged.

Electrolysis made compound water harder to abandon, without making it transparent. It turned water into an electrochemical object: something that could conduct, polarise, transport, release and exchange. The question was no longer only whether hydrogen and oxygen belonged to water. It was how electrical action could make them appear at all.

Still, electrolysis could not settle the formula. It could make water yield hydrogen and oxygen. It could not tell chemists how those elements should be counted.

That was the next difficulty. Once water had constituents, chemistry still had to learn how to count them.

References

Chang, H. (2012). Is Water H2O? Evidence, Realism and Pluralism. Dordrecht: Springer. doi: 10.1007/978-94-007-3932-1.

Davy, H. (1807). ‘The Bakerian Lecture, on Some Chemical Agencies of Electricity’. Philosophical Transactions of the Royal Society of London, 97, pp. 1–56. doi: 10.1098/rstl.1807.0001.

Faraday, M. (1834). ‘Experimental Researches in Electricity.—Seventh Series’. Philosophical Transactions of the Royal Society of London, 124, pp. 77–122. doi: 10.1098/rstl.1834.0008.

Nicholson, W. (1800). ‘Account of the New Electrical or Galvanic Apparatus of Sig. Alex. Volta, and Experiments Performed with the Same’. Journal of Natural Philosophy, Chemistry, and the Arts, 4, pp. 179–187.

Singer, G. J. (1814). Elements of Electricity and Electro-Chemistry. London: Printed for Longman, Hurst, Rees, Orme, A. Brown, and R. Triphook.

Volta, A. (1800). ‘On the Electricity Excited by the Mere Contact of Conducting Substances of Different Kinds’. Philosophical Transactions of the Royal Society of London, 90, pp. 403–431. doi: 10.1098/rstl.1800.0018.

Wilkinson, C. H. (1804). Elements of Galvanism, in Theory and Practice: With a Comprehensive View of Its History, from the First Experiments of Galvani to the Present Time. 2 vols. London: John Murray.

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It's that time of year again, when we learn which lucky scientists are among the winners of the Ig Nobel Prizes. This year, the prizes honor research on designing the perfect splash-free urinal; using mosquito proboscises to "necroprint" tiny nozzles; studying composition rates of buried cotton underwear; and the aerodynamics of a healthy nose-blow, among other highlights.

Established in 1991, the Ig Nobels are a good-natured parody of the Nobel Prizes; they honor “achievements that first make people laugh and then make them think.” The unapologetically campy awards ceremony features miniature operas, scientific demos, and "24/7 lectures," whereby experts must explain their work twice: once in 24 seconds and the second in just seven words.

Acceptance speeches are limited to 60 seconds. And as the motto implies, the research being honored might seem ridiculous at first glance, but that doesn’t mean it’s devoid of scientific merit. In the weeks following the ceremony, the winners will also give free public talks, which will be posted on the Improbable Research website.

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Arsenic contamination in drinking water poses a critical global public health hazard. Prolonged exposure to inorganic arsenic—predominantly present in groundwater as arsenate, where arsenic exists in its +5 oxidation state—can lead to severe dermatological lesions, cardiovascular diseases, and cancers. To mitigate these health risks, the World Health Organization has established a strict provisional safety guideline of 10 micrograms of total arsenic per liter of drinking water. Because arsenic occurs naturally through the weathering of rocks and minerals in aquifers worldwide, engineering low-cost, scalable, and highly efficient materials for water purification remains a central goal in modern materials chemistry.

Iron-based compounds, particularly iron oxyhydroxides, are among the most effective materials for trapping arsenate. At a fundamental chemical level, arsenate ions readily substitute for surface hydroxyl groups on iron oxyhydroxides, forming strong chemical bonds that permanently anchor the contaminant to the solid surface. However, translating this surface chemistry into practical filters presents a persistent engineering hurdle: maximizing the amount of active iron exposed to water. Simply increasing the iron concentration during material synthesis causes iron nanoparticles to aggregate into large, dense clusters. As these clusters grow, much of the iron becomes buried within the interior bulk, leaving only a tiny fraction accessible at the surface to bind arsenate.

A promising design strategy ¹ overcomes this limitation by decoupling the nucleation of iron particles from their subsequent growth on a flexible, high-surface-area polymer matrix. The foundation of this system is a composite fibrous membrane fabricated from poly(vinyl alcohol) and pectin via electrospinning. In electrospinning, a high-voltage electric field draws a liquid polymer mixture into ultra-fine threads, which solidify into an interconnected, highly porous mesh of nanofibers. Poly(vinyl alcohol) provides structural integrity and water stability upon heat treatment, while pectin—a natural polysaccharide rich in carboxyl and hydroxyl functional groups—serves as a molecular anchor capable of capturing iron ions from solution.

To grow high-density nanostructures without triggering particle agglomeration, a sequential, two-step chemical treatment is applied. First, the electrospun membrane is immersed in a solution containing iron(II) ions (Fe²⁺). The carboxylate groups on the pectin molecules coordinate with these cations, forming uniformly dispersed, nanometer-scale iron seed clusters (~2 nanometers in size) across the nanofiber surfaces as the material dries and oxidizes. In the second step, the seeded membrane is placed in a solution containing iron(III) ions (Fe³⁺). Instead of forming isolated new particles in solution, the newly added iron deposits onto the pre-existing seeds, driving controlled anisotropic growth along preferred crystallographic axes.

This seed-mediated approach yields an extraordinarily high density of surface-exposed iron nanostructures, achieving a 12-fold increase in surface atomic iron content compared to single-step doping methods. High-resolution electron microscopy reveals that these nanostructures adopt an elongated, rice-grain morphology measuring approximately 29 nanometers in width and 110 nanometers in length. Structural characterization using X-ray scattering, X-ray diffraction, and Mössbauer spectroscopy identifies the predominant active iron phase as akaganeite (β-FeOOH), a tunnel-structured, poorly crystalline iron oxyhydroxide known for its exceptional binding affinity toward arsenate.

When evaluated in water purification tests, the nanostructured membrane demonstrates exceptional superadsorbent performance. Using a minimal material dosage of 1 gram per liter, the membrane reduces initial arsenate concentrations from 1000 micrograms per liter down to below 10 micrograms per liter within 24 hours at pH 6, achieving roughly 99% arsenic removal and satisfying WHO safety limits. The material retains high efficiency across a broad operational pH range of 5 to 8, and ICP analysis confirms that iron leaching into the treated water remains safely below 5 micrograms per liter.

Kinetic and isotherm analyses reveal that arsenate removal occurs via surface chemisorption governed by cooperative binding behavior. As initial arsenate molecules bind to active sites on the akaganeite surface, local chemical and electrostatic interactions make adjacent sites even more receptive to further adsorption, resulting in accelerated uptake kinetics. Furthermore, the material performs reliably in complex water environments. Naturally occurring co-contaminants such as fluoride do not interfere with arsenate capture, and the presence of heavy metal cations—such as copper and cadmium—actually enhances arsenate removal by forming secondary surface complexes that create additional binding sites. Although chemical regeneration using sodium hydroxide solution results in a slight decrease in capacity over repeated cycles due to incomplete desorption, the membrane maintains an 81% to 87% removal efficiency, demonstrating strong potential for practical, reusable water remediation systems.

Author: César Tomé López is a science writer and the editor of Mapping Ignorance

Disclaimer: Parts of this article may have been copied verbatim or almost verbatim from the referenced research paper/s.

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When babies taste sugar for the first time, a look of delight spreads across their faces. Some wiggle with excitement. It turns out, we’re born with a sweet tooth.

“There’s a reason for that,” says Andrew Odegaard. “We crave it because our body needs it.” Odegaard works at the University of California, Irvine. As an epidemiologist, he studies health in groups of people.

Sugars, the simplest of carbohydrates, are the body’s primary energy source. Our cells run on glucose. It’s one of the main monosaccharides, the basic building blocks of sugars. Table sugar, or sucrose, is made of one molecule of glucose and one of fructose (another monosaccharide). Joined together, the two become a disaccharide (here, “di-“ means two).

Table sugar

This is a representation of a molecule of sucrose, or table sugar. It shows its two component monosaccharides. The larger molecule on the left is glucose. On the right is fructose. They are bound together at the center.

Sugar tastes sweet because its molecules bind to sweet receptors on your tongue. They signal the brain to let it know you’ve just tasted something delicious — and often full of energy.

“This role of our taste system is to act as the primary gateway controlling ingestion,” says Charles Zuker. A neuroscientist, he studies the brain at Columbia University in New York City. This system evolved, he says, “so we can be attracted to things we must have in our diet.”

Zuker is part of a team that discovered the sweet taste receptor in 2001. And this receptor requires quite a bit of sugar to make it take notice.

“It evolved that way,” Zuker says. This lets our brains tell the difference between higher- and lower-sugar foods. After all, it makes sense to pick the sweetest thing, Zuker says. “I get the most [calories] for my effort.”

In the past, when food was harder to find, that was probably very important. Today, though, many doctors worry that our diets have become too sugary.

Some people are filling up on soda and candy instead of fruits and vegetables. While some sugar is fine, eating too much can harm our health. As one way to cut back, scientists have now developed sugar-free compounds that mimic sugar’s sweetness.

Sickeningly sweet

Because we evolved to enjoy sugar, it becomes easy to eat too much of it.

Manufacturers add a lot of sugar to processed foods, such as snacks, drinks, sauces and even some breads, notes Monica Dus. She’s a neuroscientist at the University of Michigan in Ann Arbor.

When we eat more sugar than we need, our bodies will store the excess as fat. That can lead to weight gain. Over time, that extra fat can also increase someone’s risk of type 2 diabetes, obesity and heart disease, Odegaard says.

Eating sugar also causes levels of glucose to rise in the blood. When an organ called the pancreas senses these high levels, it releases a hormone called insulin.

Insulin helps glucose get into cells so they can use it for energy. But too much glucose in your blood over a long period will stress the pancreas, Odegaard says. This can lead to insulin resistance. Now, he says, “your cells can’t make use of all the sugar that you have, or they don’t work as well.”

When someone has diabetes, for example, they must be very careful about their blood glucose. That often means they need to sharply control how much sugar they eat.

No sugar added

Back in 1879, Constantin Fahlberg was a chemist at Johns Hopkins University in Baltimore, Md. One day he forgot to wash his hands when he raced out of the lab to dinner. He had been working with coal tar (a byproduct of coal). Sitting down to dinner, he was shocked to find that everything he touched tasted sweet.

He realized it was due to something he carried home from work on his hands.

Fahlberg raced back to the lab. “I tasted the contents of every beaker and evaporating dish on the table,” he later wrote. (Definitely do not try that.)

It turns out he had accidentally discovered saccharin, an artificial sweetener. Because the tongue’s sweet taste receptor is relatively insensitive, it responds to many things that aren’t sugar, Zuker explains. Some, like saccharin, bind more strongly to this receptor than sugar does.

Saccharin binds so strongly that even a tiny amount tastes intensely sweet. But our bodies don’t break down saccharin. So unlike sugar, our bodies get no energy from it.

Scientists have since developed other artificial sweeteners. The body also interprets some chemicals in fruits and other plants as sweet but doesn’t process them like sugar (see box). Some can’t be broken down by the body at all.

“You essentially just [consume] them,” he says, “and then you just pee them out.” Such sweeteners have few or no calories. They also won’t spike blood glucose.

“The biggest challenge with artificial sweeteners nowadays is that none of them really ultimately taste like sugar,” Zuker says. Some have a funny aftertaste or a strange texture. And because some (such as allulose or sugar alcohols) don’t get digested, they can make your stomach pretty unhappy.

In 2025, Zuker’s team worked out the exact structure of the sweet taste receptor. By knowing precisely how different molecules fit into it, they hope to learn how to design better low- or no-sugar sweeteners.

Too much sugar can be bad for us, but that doesn’t mean we should stop eating it, says Dus at Michigan. Glucose, after all, is the brain’s favorite dish. She says it’s unfair to say “sugar is evil” or “sugar is bad.” In fact, Dus points out, “Sugar is part of our evolutionary history.”

It’s also an important part of most cultures: Consider how widespread sweet desserts are. So just saying that sugar’s bad, she says, risks “not recognizing our biology and not recognizing our humanity.”

What is your favorite way to eat a marshmallow? Melted in hot chocolate after a long day of sledding? Or roasted and squished between a slab of chocolate and graham crackers while you are sitting around a campfire? However you like them, marshmallows are an unusual type of sweet treat — spongy, sticky and a little bit chewy, with a melting point that is just a bit above body temperature, so that they start to change from a solid to more liquid state as soon as they reach the warmth of your mouth, or the warmth of your hot chocolate or campfire!

They are also an ancient creation, originally coming from a tall marshmallow plant (Althaea officinalis) that grows in swampy fields and has a soft, spongy root. The root contains mucilage, a thick, gluey substance produced by some plants and microscopic animals to help with seed germination and water and food storage. Some cultures used the marshmallow plant to make candy, while others used it to make medicine. The ancient Egyptians, for example, dried the root and mixed it with honey to make marshmallow treats, but the early French thought the root looked like lung tissue, and experimented with making medicine out of its gummy juice to soothe sore throats.

Today’s modern marshmallows contain no parts of the marshmallow plant though, because the plant is considered a medicinal herb. Instead, modern marshmallows are primarily a mix of three ingredients: sugar, corn syrup and gelatin, with the gelatin replacing the thick, gluey substance from the marshmallow plant.

The sugar and corn syrup are heated together with a little water to make a solution. The solution is warmed and brought to a boil, turning it into syrup. The sugar in the solution makes the boiling point of the solution higher than the boiling point of pure liquid water, which is 100° Celsius (212° Fahrenheit). The greater the concentration of sugar in a solution, the higher its boiling point. As the solution boils, water evaporates, the solution becomes even more concentrated with sugar and the boiling point gets pushed even higher. For example, an 85 percent sugar solution has a boiling point about 16 degrees C (28 degrees F) higher than the boiling point of pure water.

The final concentration of sugar in the syrup determines the structure of the candy that forms when the syrup cools. Think about the difference between the texture of caramels and lollipops — the caramel is softer and chewier, while the lollipop is hard and cracks when it is bitten. The syrup used to make caramels is cooked until it reaches 115–120 °C (240–250 °F), at which point it has an 87 percent concentration of sugar in solution and enters the firm ball stage (see the Exploratorium link in the Bibliography for more information about this). The syrup used to make lollipops, though, is cooked until it reaches a 99 percent concentration of sugar in solution at a temperature of 150–155 °C (300–310 °F), which makes drops of the syrup turn into hard, brittle, easy-to-break threads when placed in cold water.

When making marshmallows, the syrup is cooked until it reaches 112–115 °C (235–240 °F) or 85 percent concentration of sugar in solution, and corn syrup is added to the solution to help prevent crystals from forming in the cooled syrup. Nobody likes crunchy marshmallows!

So, where does the gelatin come in, and just what is gelatin anyway? Gelatin is that wiggly, jiggly food that seems almost magical as it suspends fruits, forms shapes and sculptures and transforms from a solid to a liquid as soon as it hits the warmth of your mouth. It is not magic, but special gelatin molecules are responsible for these food tricks.

Gelatin comes from collagen, the main protein in connective tissue, which is a specialized, fibrous tissue in animals that connects other tissues together, like tendons connect muscles to bone. The special quality of gelatin is that it is a protein that can coagulate (or come together) when it is beaten or whipped. So, to make marshmallows, hot concentrated syrup is beaten into gelatin with a mixer. The beating process forms bubbles in the syrup and protein mixture, and the gelatin coagulates at the point where the air in the bubble and the syrup-protein mixture meet, stabilizing the bubble walls so they don’t collapse.

In this cooking and food science fair project, you will make your own gooey marshmallows, using different amounts of sugar and corn syrup to find out which ratio of sugar to corn syrup makes the best-tasting and best-textured treat.

Experimental procedure

Preparing to cook

1. Wash your hands.
2. Get out all cooking tools and ingredients and set them out so they are ready to go and easy to access.
3. Identify the 116 °C mark (240 °F) on your candy thermometer. This is the traditional temperature used to make marshmallows.

Preparing your cake pans

1. Label the bottoms of three cake pans with masking tape pieces, numbered 1 to 3. Pan 1 will contain recipe 1, pan 2 will contain recipe 2, and pan 3 will contain recipe 3. Each recipe contains a different ratio of sugar to corn syrup.
2. Pour a small amount of vegetable oil on a paper towel and lightly oil the three cake pans.
3. Pour a small amount of powdered (confectioner’s) sugar in the strainer and dust the inside of the cake pans.
   1. If you get too much powdered sugar in the pan, hold the pan upside-down over a sink, and gently tap out the excess.

Preparing your gelatin

1. Pour 40 mL (1/6 cup) of cold water into the large mixing bowl.
   1. To get 1/6 cup, fill the 1/3 measuring cup halfway with water.
2. Sprinkle one 1/4-ounce (oz) envelope of plain, unflavored gelatin over the cold water.
3. Mix the gelatin and water together for about 5 seconds (sec) with a fork and set the bowl aside. This will give the gelatin time to “soften” or “bloom,” as shown in Figure 1 below.

Preparing your syrup

1. Add 60 mL (1/4 cup) of cold water to the saucepan.
2. Using Table 1 below as a guide, add granulated sugar and corn syrup to the water in the saucepan, according to which recipe you are currently making. This step is shown in Figure 2 below.

3. Put the lid on the saucepan. An adult should closely supervise all work from this point on.
4. Turn on the stove to medium-high heat.
5. Lift the lid and check the solution in the saucepan about every 30 sec until it just comes to a boil.
6. Remove the lid and set it aside.
7. Begin to measure the temperature of the syrup in the saucepan using the candy thermometer.
   1. When using the thermometer, do not let the tip of the thermometer touch the bottom or sides of the pan. Instead, try to put the tip below the surface, closer to the middle of the pan.
8. Continue heating your syrup solution until the temperature reaches 116 °C (240 °F).
   1. WARNING: The syrup will be very hot, so use extreme caution when working around the pan.
   2. Observe how the temperature rises. Does it go up quickly at first? What happens as the solution becomes more concentrated? Record your observations in your lab notebook.
9. When the temperature reaches 116 °C (240 °F), immediately turn off the stove and move to the next section, Making Your Marshmallows.

Making your marshmallows

1. Turn the mixer on low and carefully pour the hot syrup solution into the gelatin and water in the large mixing bowl, as shown in Figure 3 below.
   1. Note: Some syrup solution will likely solidify on the saucepan as the syrup cools. However, if you pour too slowly, or wait too long and do not immediately pour the syrup into the bowl after turning off the stove, you may end up with a lot of syrup solidified on the saucepan, which may affect your results.

2. Set the timer for 11 minutes and start it.
3. Gradually increase the speed of the mixer until it is operating at full (high) speed, as shown in Figure 4 below.

4. Continue to beat for approximately 11 minutes, or until the mixture starts to become very thick, glossy and lukewarm.
5. Add 2.5 mL (1/2 teaspoon) of vanilla, as shown in Figure 5 below, and beat for another minute.
   1. Write down the total time that you beat the recipe in your lab notebook.
   2. Try to beat each recipe for approximately the same amount of time.

6. Pour a thin layer of vegetable oil on a spatula and scoop out the marshmallow mixture from the mixing bowl and into the prepared cake pan with the masking tape piece that matches the recipe number.
7. Use the spatula to gently smooth down the top of the marshmallow, so that it is pretty flat.
   1. If the top of the marshmallow mixture naturally smooths itself down when you put it in the pan, or if the mixture is so sticky that it is very difficult to smooth it down with the oiled spatula, skip this step.
8. Wash and dry all cooking utensils.
9. Repeat the sections Preparing Your Gelatin, Preparing Your Syrup and Making Your Marshmallows two more times with clean materials, until all three recipes have been prepared, as shown in Figure 6 below.
   1. To have the best control over all the variables you should make all three recipes on the same day. Otherwise, uncontrollable variables (like the humidity on that day) may affect your results.
10. Repeat the entire Experimental Procedure (up to this point) two more times so that you have a total of three trials (pans) for each recipe. Repeating the recipes makes sure that your test results are accurate and repeatable.
    1. You can do this step on a different day if you would like.

11. Allow the three marshmallow “pies” to sit out, uncovered, on a counter for at least 4 hours, or up to overnight, so that they become firm.

Preparing your marshmallows for testing

1. Once the marshmallow pies have become firm, after at least 4 hours or overnight, turn the pans over, one at a time, onto a cutting board.
   1. If the marshmallow sticks to the bottom of the pan, push on the bottom of the pan a bit, or use a spatula to help lift it out.
2. Roll a pizza cutter in some powdered sugar, and then use it to cut up your marshmallow into pieces that are about 3 centimeters (cm) by 3 cm (one inch by one inch). Use a ruler to help measure where to cut.
   1. An alternative to the pizza wheel is to use a small cookie cutter to cut out shapes.
3. Dust the marshmallows on all sides with a little powdered sugar, using the strainer.
4. Place the marshmallows in the air-tight container or plastic bag, being sure to label the container or bag with the recipe number and the trial number (remember, you will be performing three trials) from the masking tape pieces, so you know which recipe was used for each bag or container. The marshmallows will keep in an airtight container for up to a week, so you have around six days in which to make and test all the trials.

Testing your marshmallows

Perform the following three tests to investigate the melting temperature, the toughness and the taste of the marshmallows you made from each recipe. Perform each test one trial set at a time.

1. Melting tests: Drop three marshmallow pieces (one from each recipe) in a pan of hot water and see which piece melts first.
   1. Because of slight variations in marshmallow sizes, pick three pieces that are of the same size and shape.
   2. Start a timer and see which one melts first, second and last. Record your results in a data table in your lab notebook.
   3. Repeat the test at least three times per trial set to make sure your results are repeatable and accurate.
2. Cutting tests: See which marshmallow is the toughest using a cheese cutter, like the test described in the Science Buddies cooking and food science fair Project Idea Tough Beans (in the Procedure section titled “Testing Your Cooked Beans”). Alternatively, you can try an easier but less quantitative method using a small disposable cup, like a Dixie cup.
   1. If you use the disposable cup method, do the following:
      1. Put a marshmallow piece from one of the recipes on a cutting board.
      2. Arrange a small cookie cutter so that it can sit on top of the marshmallow. (You will need the cookie cutter to be thinner than the marshmallow pieces to do this.)
      3. On top of the cookie cutter, balance the small disposable cup.
      4. Put 20 pennies in the cup, one at a time.
      5. Remove the cup and cookie cutter from the marshmallow piece and inspect the marshmallow. Did the cookie cutter leave a noticeable impression? Use a camera to take a picture of the marshmallow piece.
      6. Repeat this process with marshmallow pieces from the other recipes. Eventually you will want to have repeated the test at least three times per trial set. Record your results in a data table in your lab notebook.
3. Taste tests: Gather your family and friends and give them three samples of marshmallows, one from each recipe in a trial set. Have them evaluate the samples for taste and texture.
   1. Which recipe do they think made the best-tasting samples? Which recipe do they think made the best-textured samples? Record your results in a data table in your lab notebook.
   2. Read this Science Buddies article about sample size to determine how many volunteers you will need.

Analyzing your results

1. For melting tests, average the results of your time tests from the three trials, and plot the decimal ratio of sugar to corn syrup on the x-axis, and the melting time on the y-axis.
   1. Do you see a correlation between the melting time and the ratio of sugar to corn syrup used? If so, why do you think this is? Hint: Re-read the introduction.
2. For cutting tests, if you used the protocol from the science project Tough Beans, average your coin counts from the three trials (to get a measure of the force required to cut through the marshmallow), and plot the decimal ratio of sugar to corn syrup on the x-axis, and coin counts (or cutting force) on the y-axis.
   1. If you used the small disposable cup method, print out the pictures you took and put them on your science project display board.
   2. Do you see a correlation between how tough the marshmallow pieces were and the ratio of sugar to corn syrup used? If so, why do you think this is?
3. For taste tests, make a bar chart showing the ratio of sugar to corn syrup on the x-axis, and the number of people who thought it had the best taste on the y-axis. Make a second bar chart showing the decimal ratio of sugar to corn syrup on the x-axis, and the number of people who thought it had the best texture on the y-axis.
   1. Why do you think some people preferred marshmallows made using one recipe compared to marshmallows made using another recipe? Were some pieces too sweet or too hard for your taste testers, or not sweet enough or too soft?

Variations

• Develop other ways to test your marshmallows. For example, measure the density of the marshmallows from each recipe, using a scale, graduated cylinder and Archimedes’ principle. Plot the ratio of the sugar to corn syrup on the x-axis and the density on the y-axis.
• Choose the recipe that has a decimal ratio of 2.7 for the sugar to corn syrup. Vary the final cooking temperature (deviate from the traditional 116 °C [240 °F]) to see how that affects the taste, texture, melting time, toughness and/or density. Be sure to conduct repeat trials at each temperature you test.
• Break up your taste-testing volunteers into male and female, or by age, to see if there are gender or age differences in marshmallow taste or texture preferences.

This activity is brought to you in partnership with Science Buddies. Find the original activity on the Science Buddies website.

Two are dead and over 25 buildings damaged after a cataclysmic explosion at a cannabis facility in Sand Springs, Oklahoma last week.

According to initial reporting by KTUL, the blast happened at 1:30am, and literally blew the roof off the building, sending out shockwaves that could be felt up to 30 to 40 miles away.

“It was like around 1:20, 1:25. And it was just a big boom,” local man Donovan Ruffin, who lives a mile away from the blast, told the news channel. “Me and my wife woke up at the same time. And as soon as I heard it, I woke up… I have a window right by my bed, so I seen a flash.”

Ryan Bassham, the owner of the THC facility where the explosion occurred, was later arrested on murder charges, People reported, after defying an order from the fire marshal to halt extraction operations.

One explanation being offered up — though not officially confirmed — is that the explosion traces back to the use of butane to produce THC concentrate.

One of the more popular methods by which the cannabis industry produces concentrated products like rosin is through a process known as butane hash oil (BHO) extraction. Basically, producers use liquid butane as a solvent, pouring it over raw plants at temperatures in the ballpark of -40 degrees Fahrenheit.

Because liquid butane has an exceptionally low boiling point — 30.02 degrees Fahrenheit to be exact — the solvent can be easily cooked off after mixing with the plant matter, leaving just the dissolved cannabinoids behind. Though this is said to be one of the most straightforward methods for concentrating cannabis, the use of highly flammable butane makes it one of the most dangerous.

As the publisher Technology Networks explained in a blog post on BHO extraction, a “faulty gasket or improper check could lead to disaster.”

Overall, there is no national count of BHO-related explosions or injuries, so it’s difficult to say how common an occurrence it is. According to one clinical study conducted at the burn unit of the University of California, Davis Medical Center, there were at least eight BHO-related hospitalizations between 2011 and 2012, though that number is likely higher today after a decade of legalization and cannabis industry growth.

More on THC: Research Suggests the Older You Get, the More Weed You Should Smoke

The post Marijuana Factory Explodes, Causing Two Deaths appeared first on Futurism.

Author: José Luis Granados Mateo is a postdoctoral researcher in the Department of Philosophy at the University of the Basque Country (EHU) and a member of the Integrated History and Philosophy of Science (iHPS) research group. His work focuses on history and philosophy of science, science and values, and the epistemology of scientific practices.

Take a glass of water. Nothing about it seems to ask for a theory. It has no colour worth naming, no dramatic smell, no visible architecture. It runs, freezes, boils, evaporates, dissolves, erodes, circulates and nourishes. It belongs to weather, geology, physiology, cooking and plumbing before it belongs to chemistry.

Then science gives us the famous abbreviation: H₂O.

There is a curious finality in those three symbols. They seem to do what ordinary experience cannot. Rain, steam, ice, rivers, blood and laboratory wash bottles are all brought under a clean chemical statement. Two atoms of hydrogen, one atom of oxygen. For many people, H₂O is more than a formula. It is almost an emblem of scientific knowledge, the moment when science gets behind appearances and tells us what something really is.

Yet the formula stands at the end of a long story. Before water could be written as H₂O, it had to acquire a chemical interior. Chemists first had to learn to treat the most familiar liquid in the world as something that could be taken apart, recomposed and assigned constituents. The first transformation was not from one formula to another. It was from element to compound.

This is the starting point of Hasok Chang’s Is Water H₂O?, a book that does for water something close to what Inventing Temperature did for thermometry. Temperature now looks like a number waiting to be read from a scale. Chang showed that the scale had to earn that appearance of simplicity. Fixed points had to be stabilised, instruments compared, extrapolations justified. Measurement became routine only after it had first been constructed as a reliable practice.

Water presented another kind of difficulty. The problem was not how to make a phenomenon answer numerically, but how to make a familiar substance answer analytically. Chemists had to learn to look at water as more than something that flowed, froze, boiled, dissolved, condensed or could be purified. They had to ask whether it had an inside: whether it could be taken apart, made again, and understood through its constituents.

To us, that question feels almost unavoidable. In the eighteenth century, it was anything but.

When water had no parts

For much of European natural philosophy, water was an element. That claim can sound childish to us, though only because we hear it through the categories that later displaced it.

The idea was older than Aristotle. Thales had imagined water as the origin of things; Empedocles later placed it among the four basic “roots” of the world, alongside earth, air and fire. Aristotle gave that inherited scheme a more systematic form. But these were not elements in the modern chemical sense. No one was classifying them by atomic number, nuclear charge or molecular structure. They belonged to a physics of qualities and transformations, of hot and cold, wet and dry. Water was simple because, in that world, it occupied a basic place in the order of natural change.

By the eighteenth century, that older grammar had been heavily rewritten. Chemistry had become crowded with acids, alkalis, precipitates, metallic calxes, distillations and newly isolated “airs”. Water was everywhere in this world. It served as solvent, medium, reagent, product and experimental environment. It could be purified, frozen, boiled, condensed and recovered.

Still, none of that made it automatically compound. A substance can be experimentally familiar and still have no recognised chemical interior.

The difficulty is easy to miss because we already speak the language that solved it. We say that hydrogen burns in oxygen to form water. But eighteenth-century chemists did not begin with hydrogen and oxygen. They had inflammable air, dephlogisticated air, fixed air, nitrous air and other gaseous bodies whose identities were still unsettled. The experiment did not come with modern subtitles. The vocabulary in which its meaning now appears self-evident was itself one of the products of the revolution.

Water became a compound when certain operations began to count as analysis and synthesis: taking water apart, producing it again, and treating the substances involved as its constituents. This was more than a taxonomic adjustment. It changed the standard of chemical knowledge. To know a substance now meant, increasingly, to know what it was made of.

The wrong theory that worked

The defeated rival was phlogiston theory. Posterity has enjoyed defeating it again and again.

In the usual caricature, phlogiston was a ghostly fire-stuff. Combustible bodies contained it and released it when they burned. Metals contained it too. When heated in air, they lost phlogiston and became calxes. Heat the calx with charcoal, and the metal returned because phlogiston had been restored. Then Lavoisier introduced oxygen, and the ghost vanished.

What makes this story so satisfying? Precisely what makes it misleading. It tells the Chemical Revolution from the winner’s seat, where oxygen looks inevitable and phlogiston appears only to make its exit.

From inside the eighteenth-century laboratory, the story looks less like a comedy of error. Phlogiston was not a fantasy floating above the experiments. It belonged to things chemists and metallurgists handled every day. Metals dulled and lost their malleability when calcined. Calxes could be brought back to metal with charcoal. Ores became usable metals in furnaces. Fire, air, metals, respiration and reduction seemed to be tied together by a common traffic of activity. Phlogiston named that traffic. It was a mistaken name for a real pattern.

No one shows this better than Joseph Priestley. He was not a dim survivor of pre-modern chemistry, nor an experimental genius trapped in mere confusion, but one of the great investigators of the eighteenth century, a dissenting minister, political radical and virtuoso of pneumatic chemistry. In Leeds, living beside a brewery, he studied the “fixed air” rising from fermentation, now carbon dioxide. Later, by heating the red calx of mercury, he obtained a far more remarkable gas. Flames burned in it with unusual brilliance; animals could breathe it for longer than ordinary air; Priestley even tried it himself, noticing what he described as a peculiar lightness in his chest.

We call that gas oxygen. Priestley did not. To him, it was dephlogisticated air.

The name sounds absurd only when detached from the operations that made it meaningful. In phlogiston chemistry, ordinary air could become loaded with phlogiston. Air deprived of phlogiston would therefore be especially ready to absorb it again. It should support combustion powerfully and sustain respiration unusually well. That is exactly what Priestley saw. His interpretation was wrong, but it was not arbitrary. It made sense of real experimental behaviour.

Nor was the winning term free of theory. Lavoisier called the gas oxygen, meaning “acid-generator”, because he believed oxygen was the universal principle of acidity. That too was false. The defeated word preserved one mistake; the victorious word preserved another. One became a fossil of discarded chemistry. The other became familiar enough to hide its own history.

A few drops on the glass

The next turn in the story came from a gas with a wonderfully plain eighteenth-century name: inflammable air. We now call it hydrogen.

Henry Cavendish had studied this light, combustible gas in the 1760s. It could be produced by reacting metals with acids, and it burned readily. When inflammable air was burned together with dephlogisticated air, something extraordinary happened. Water appeared. Two invisible airs had produced the most familiar liquid in the world.

It is hard, now, not to let modern chemistry finish the sentence for us. Hydrogen combines with oxygen to form water. Of course. But the droplets on the glass did not settle the matter by themselves. They had to be interpreted.

A phlogistonist could tell the story differently. Perhaps inflammable air was water combined with phlogiston. Perhaps dephlogisticated air was water deprived of it, or at least air peculiarly ready to receive it. On that reading, the experiment did not reveal that water was composed of two gases. It showed that ordinary water could be restored through the right exchange of phlogiston.

That may sound evasive from our side of the Chemical Revolution. At the time, it was not an absurd reply. A flame, a gas, a misted vessel, a change in weight. None of these comes with its meaning attached. Evidence becomes evidence inside a working system of concepts, operations and standards. What counts as a substance? What counts as a component? What counts as analysis rather than transformation? Which variable should be trusted most, appearance, activity, weight, or reproducibility?

The new chemistry changed the setting in which the experiment was read. The appearance of water was no longer merely a striking event at the end of combustion. It became part of a quantitative account. How much inflammable air had disappeared? How much oxygen had been involved? How much water had been produced? Could water be decomposed and then recomposed, with the numbers matching closely enough to make the same substance appear on both sides of the operation?

At that point, water was no longer simply appearing on the glass. It was entering the balance sheet of chemistry.

The balance makes composition visible

A metal is heated in air. It dulls, loses its shine, gives up its malleability and turns into a calx. In phlogiston chemistry, this looked like a loss. The metal had given up the principle that made it metallic.

Then came the awkward fact on the bench. The calx was heavier.

Heavier, after a supposed loss.

That did not destroy phlogiston theory in a single blow. Good chemists could still defend it, and did. Yet the weight gain gave oxygen chemistry an opening. Perhaps calcination was better understood as combination with something from the air. Follow the weights carefully and chemical change begins to look different. Appearances still matter. Activity still matters. So do colour, heat, flame, residue and reversibility. But weight gives the whole operation a new discipline. It turns transformation into material accounting.

The same habit of attention changed the status of water. A substance that can be broken down into measured quantities of other substances, and then produced again from them, starts to lose its old simplicity. It has become analysable. It has constituents.

This was the strongest hand of the new chemistry. Combustion, calcination, gases and water could now be drawn into the same quantitative order. Chemical names changed too. A name was expected to do more than preserve origin, appearance, use or inherited convention. Ideally, it should tell chemists what a substance contained.

The point needs care. This was never a struggle between careful measurers and careless speculators. Cavendish was one of the most exact experimentalists of the century. Richard Kirwan, among the serious defenders of phlogiston, treated weight as a genuine difficulty. The disagreement concerned the authority of weight. Should it count as one experimental constraint among others, or should it become the privileged route to chemical identity?

That question slowly shifted the terrain. Phlogiston chemistry was most at home in a language of principles, powers and transformations. What gives a body combustibility? Why does a metal lose its metallic character? How can a calx be revived? What passes from one substance to another when reaction occurs?

The new chemistry pressed elsewhere. What does this substance contain? In what proportions? Can those constituents be separated and brought together again while the material account remains intact?

Once that style of reasoning became dominant, water’s old simplicity became harder to defend. Water could still be boiled, condensed, purified and recovered, just as before. Now it could also be weighed into a different story. It could be decomposed, recomposed, renamed and treated as a body with an interior.

No single knockout blow

The usual story of the Chemical Revolution likes a clean scene. Priestley isolates oxygen and fails to recognise it. Lavoisier understands what Priestley could not. Phlogiston collapses. Modern chemistry begins.

History was less obliging.

Priestley knew the new chemistry and resisted it. Cavendish remained cautious. Other competent chemists hesitated, adopted parts of the oxygen system, or kept some version of phlogiston alive. Their reluctance was not simple blindness. Phlogiston theory still connected combustion, respiration, calcination, reduction and the behaviour of airs. It had also helped produce many of the experimental facts that later counted in favour of oxygen chemistry.

Lavoisier’s system, for all its power, carried its own burdens. His theory of acidity was wrong. The very word oxygen preserved that mistake, since it named the gas as the supposed generator of acids. His chemistry also relied on caloric, an imponderable fluid of heat. Seen from the present, caloric does not look obviously more respectable than phlogiston. Lavoisierian chemistry was elegant, disciplined and transformative. It was not modern chemistry arriving whole and immaculate.

So the real historical question is more interesting than the textbook one. Why did one way of doing chemistry become more convincing, more portable, more teachable and more productive than another?

Much of the answer lies in the rise of compositionism. Oxygen chemistry made substances legible as combinations of measurable constituents. Reactions became material accounts. Chemical names began to function as compressed statements of composition. Laboratory operations, weights, gases and nomenclature could now be made to speak the same language.

Phlogiston still made sense of many transformations, but it sat uneasily in that emerging world. It named a principle of activity, a way of thinking about combustibility, metallic character, reduction and chemical power. The new chemistry increasingly wanted components. It asked what substances were made of, how much of each constituent they contained, and whether those constituents could be separated and brought together again.

As compositionism gained ground, phlogiston began to lose its natural habitat. It could still speak about combustion, reduction and chemical activity, but chemistry was moving towards another language: constituents, proportions, balances, material accounts. The revolution changed the answer, certainly. More importantly, it changed the kind of answer chemistry wanted.

The first metamorphosis of water

When water became a compound, it did more than shift from one scientific category to another. Its identity for chemistry changed.

Before this transition, water could pass through laboratory practice while keeping its apparent simplicity. It could be boiled, frozen, distilled, condensed, purified and recovered. It appeared everywhere, as liquid, vapour, solvent, medium, residue or product. Familiarity gave it stability. It did not yet give it an interior.

After the Chemical Revolution, the same substance began to occupy a different role. Water could be produced from other bodies, analysed through them, and assigned constituents. The liquid that returned from vapour or appeared after combustion now carried a hidden depth. It was no longer merely recovered; it could be accounted for.

Nobody saw that hidden constitution. There were no water molecules under a microscope, no direct image of hydrogen and oxygen joined together. The evidence came through laboratory work: gases prepared and collected, vessels sealed, mixtures burned, droplets condensed, weights compared, names revised. Compound water became credible through practice before it became imaginable as a molecular structure.

This is why the story should not run too quickly towards H₂O. Lavoisierian chemistry helped establish water as a compound of hydrogen and oxygen. It did not settle the formula. It was one thing to say that water was made from hydrogen and oxygen; it was another to decide how many atoms of each entered the compound. That question would take nineteenth-century chemistry into a different set of difficulties.

First came the more basic metamorphosis. Water ceased to be one of nature’s simple terms. It became a substance that could be decomposed, recomposed and written in the language of other substances.

The cost of clarity

The new chemistry clarified immensely. It disciplined measurement, stabilised nomenclature and helped make possible the modern science of composition. But clarity selects. It illuminates some things by pushing others into shadow.

Phlogiston theory had been concerned with chemical activity. Why do substances burn, reduce, revive, calcine, combine and transform? Early compositional chemistry was much better at saying what entered and left a reaction than at explaining why the reaction happened at all. It provided an excellent account book, though not yet a full theory of chemical agency.

Some questions once handled by phlogiston later returned in different forms. Chemical potential energy, oxidation-reduction and electrons all came to occupy parts of the conceptual space that phlogiston had once held. That does not vindicate phlogiston as a literal substance. It suggests something subtler and more useful. A false theory can preserve real questions. A victorious theory can be right and still leave work unfinished.

This is a more realistic picture of scientific progress. Oxygen chemistry won for good reasons. But those reasons were not contained in one devastating experiment before which only fools could hesitate. They belonged to a wider reorganisation of chemical practice.

Water’s first great chemical transformation was therefore not the discovery of a formula. It was the making of a new kind of object. The familiar liquid became something that could be analysed, recomposed, weighed and named through its constituents. Only after that could the later question become pressing: not whether water was compound, but what exactly its composition was.

Electricity seemed about to settle the matter. It made water yield hydrogen and oxygen directly, as if nature had finally supplied the decisive experiment.

Instead, the proof produced a new puzzle.

References

Chang, H. (2012). Is Water H₂O? Evidence, Realism and Pluralism. Springer. doi 10.1007/978-94-007-3932-1.

Chang, H. (2004). Inventing Temperature. Measurement and Scientific Progress. Oxford University Press.

Cavendish, H. (1784). ‘Experiments on Air’. Philosophical Transactions of the Royal Society of London, 74, 119–153.

Kirwan, R. (1789). An Essay on Phlogiston and the Constitution of Acids. London, J. Johnson.

Lavoisier, A.-L. ([1789] 1965). Elements of Chemistry. Dover.

Priestley, J. ([1796] 1969). Considerations on the Doctrine of Phlogiston, and the Decomposition of Water. Kraus Reprint.

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Fashionistas of a certain age no doubt remember the late designer Stephen Sprouse, who gave Blondie's Debbie Harry her distinctive Bowery punk style. Sprouse's first collections in the mid-1980s made a major splash, melding graffiti prints, a playful '60s aesthetic, and bright Day-Glo colors that won admirers such as Andy Warhol.

Several examples of Sprouse's work are preserved in the textiles collection of the Indianapolis Museum of Art at Newfields, where conservationists are trying to learn more about the phosphorescent and fluorescent pigments the designer used in his creations in order to keep the colors from dulling over time. They described recent advances in that research at a meeting of the American Chemical Society (ACS) held this week in Chicago.

Scientists began describing minerals that glowed in the dark as phosphors in the Middle Ages. Typically, such materials take on a charge when exposed to light, storing the absorbed energy and releasing it gradually as re-emitted light. Henri Becquerel discovered radioactive decay in 1896 because of his work with phosphorescent uranium salts: He left the salts in a closed drawer with photographic plates, which fogged up even without a light source. DayGlo paints and similar materials exploit the related phenomenon of fluorescence. DayGlo pigments fluoresce in daylight by essentially converting UV light into visible light to produce brighter, more vibrant colors.

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2,000 years ago, ancient Roman engineers were building structures that have withstood the test of time remarkably well, many of which remain standing today. At the heart of their success lies a mystery: what were the secrets of ancient Roman concrete that made these buildings so durable?

Now, MIT researchers are drawing from this curiously durable technology of the ancient world to help them develop modern building materials that can last longer, and potentially even reduce the carbon footprint of modern construction—all while managing to self-repair itself by healing small cracks.

Reverse-Engineering Roman Concrete

The mystery of Roman concrete became a fascination for MIT Associate Professor Admir Masic, who has spent many years now investigating the chemistry and other properties of this durable ancient technology. Based on his research, the material’s longevity isn’t just a product of the ingredients used by its ancient creators: the secret of Roman concrete also has to do with its manufacturing process.

This, according to Masic, seems to be at the heart of the material’s unusual ability to repair itself after cracks begin to form.

In 2023, Masic and his colleagues began looking at samples of lime that were very rich in calcium, which researchers call “lime clasts.” These clasts are ubiquitous throughout concrete found in ancient Roman constructions, although they were long discounted as being the byproducts of imperfections in the mixing process.

However, closer inspection revealed something remarkable about these lime clasts: as cracks form in the concrete, water can gradually enter the material, allowing calcium within the lime casts to dissolve and then recrystallize, effectively sealing the newly formed openings.

Applying Ancient Technology to Modern Engineering

Now, that discovery is propelling the development of modern concrete that incorporates calcium-rich lime clasts, thanks to a partnership between Masic and Italian entrepreneur Paolo Sabatini, who joined forces to develop a commercial version now being produced by Dmat, a company they co-founded in 2021.

Part of the strength of Dmat’s approach is the fact that they have developed additives that can be added to existing kinds of concrete and mortar, rather than requiring builders to acquire new and unfamiliar materials.

“We can now offer an extremely competitively priced, self-healing product that is easy to implement and available worldwide,” Masic recently told MIT News.

“What’s exciting to me is that this material could become the industry standard without requiring companies to change how they operate,” he added.

Increased Structural Lifespan and Environmentally Friendly

Beyond its ability to strengthen the durability of concrete, Dmat’s new technology can also reduce carbon dioxide emissions by up to 40 percent compared to those in traditional concrete. Presently, the new material is already being used in several infrastructure projects in parts of Europe, with eventual plans for expansion into the United States and other countries.

Widespread implementation and scalability like this is important, according to Masic and his colleagues, since large volumes of concrete are produced throughout the world, the production of which accounts for as much as 8% of global carbon dioxide emissions.

Extending the life of structures while simultaneously reducing the emissions associated with their construction would therefore hold significant environmental benefits.

Building the Future with Lessons from the Past

“At Dmat, we like to view concrete as an ecosystem,” Sabatini recently said in a statement. Successfully introducing a new material, he argues, requires considering not merely its chemistry, but cost, transportation, certification, and compatibility with established construction practices.

Notably, Dmat’s approach differs from other examples of modern experimental self-healing concretes, which generally rely on bacteria or polymers. Instead, by placing the focus on ancient techniques used by Roman engineers that long remained overlooked and misunderstood, Masic and his colleagues have developed an innovative new technology with roots that go back almost two millennia.

“There’s something profound about how ancient builders, without our modern chemistry, engineered self-healing material that still stands today,” Masic notes.

“My group research and work with Dmat is to make the modern built environment better by applying the best lessons from the past to today’s challenges,” he adds.

Micah Hanks is the Editor-in-Chief and Co-Founder of The Debrief. A longtime reporter on science, defense, and technology with a focus on space and astronomy, he can be reached at micah@thedebrief.org. Follow him on X @MicahHanks, and at micahhanks.com.

One of the biggest challenges in the search for alien life is that Earth remains the only example of life we know. As a result, nearly every attempt to identify biology elsewhere starts, at least partly, by looking for chemistry that resembles life here.

Researchers at the University of Glasgow are exploring a different strategy. Rather than asking whether an alien molecule looks biological by Earth standards, they propose asking a more fundamental question: How difficult was that molecule to make?

In a new study published in the Proceedings of the National Academy of Sciences, researchers demonstrate a method that combines mass spectrometry with machine learning to estimate what researchers call “molecular assembly,” or MA.

The concept comes from Assembly Theory, which attempts to quantify molecular complexity by identifying the minimum number of construction steps needed to build a molecule. Researchers argue that when highly complex molecules appear repeatedly and in abundance, their existence may point to a process that is consistently steering chemistry down particular pathways.

That could make molecular assembly useful as an “agnostic” biosignature—one that does not depend on alien life sharing Earth’s chemistry.

“There is a growing interest in ‘agnostic’ life detection strategies: methods that can identify molecular signs of life without needing an exact understanding of what that life might look like, or how it might interact with a complex, poorly characterized environment,” the researchers write. “The measured assembly index, defined as the number of steps along the shortest construction pathway, is fixed for each molecule and does not depend on external conditions, such as the environment in which the molecule might be detected or the specific ‘alien’ biochemistry from which it might arise.”

The researchers’ proposed “agnostic” approach could become increasingly important as scientists encounter tantalizing, but often ambiguous, chemical signals in the search for alien life.

Methane on Mars, phosphine reported in Venus’s atmosphere, and dimethyl sulfide proposed as a possible biosignature within exoplanet atmospheres have all attracted attention because living organisms can produce those compounds.

However, biology is not the only way to make them. Methane, for example, can emerge from geologic processes. Dimethyl sulfide is strongly associated with life on Earth, yet it has also been detected in cometary material and produced in laboratory experiments simulating primordial chemistry.

The presence of a particular molecule, in other words, does not necessarily tell scientists what produced it. Molecular assembly approaches the problem from a different angle. Under Assembly Theory, researchers calculate the shortest possible pathway for constructing a molecule from simpler building blocks. A molecule with an MA of 40, for example, has a shortest construction pathway twice as long as one with an MA of 20.

The reasoning is rooted in probability. Undirected chemistry has an enormous range of possible molecular structures. As molecules become more complex, the odds of repeatedly producing the same highly elaborate structure shrink unless something constrains the process.

Biological systems store and use information to direct chemical reactions toward particular outcomes, permitting organisms to manufacture the same complicated molecules again and again instead of wandering randomly through the immense space of possible chemistry.

Previous experiments cited in the study found an approximate dividing line between abiotic and biological samples at an assembly index of around 15.

Perhaps more importantly, in the search for alien life, scientists would not necessarily need to know what a molecule is before recognizing that it is unusually complex.

That could be especially useful on worlds such as Mars, Europa, Enceladus, or Titan, where spacecraft may encounter compounds unlike anything commonly found on Earth.

Several proposed planetary missions use, or could use, mass spectrometers. These instruments analyze molecules by measuring ions’ mass-to-charge ratios.

More sophisticated multistage mass spectrometers can repeatedly break molecules apart, giving indications of their structure and assembly history. But spacecraft operate under severe limits on power, mass, computing capacity, and instrument complexity. Many planetary missions, therefore, rely on simpler, single-stage mass spectrometry.

To close that gap, the researchers trained an XGBoost model on standardized mass spectra from the National Institute of Standards and Technology Chemistry WebBook. The model learned to estimate molecular assembly directly from single-stage spectra.

According to the study, it reduced relative mean squared error threefold compared with the best-performing baseline model, cutting the figure from 0.12 to 0.04.

The model also tended to underestimate molecules with especially high MA values. For a life-detection system, that bias could be preferable. It means the model was more likely to be conservative than to classify a molecule as an unusually strong biosignature mistakenly.

Researchers envision future spacecraft using this kind of analysis autonomously. A rover or lander that detects molecules with unusually high assembly scores might flag the surrounding area for closer inspection, prioritize certain samples, or identify material worth returning to Earth.

The study explicitly evaluated conditions designed to resemble instruments proposed for missions to Titan and Europa.

The same framework also raised another, more speculative possibility: when the researchers compared very large molecular databases, they identified several medical compounds with unusually high MA values. Their complexity was not the product of biology alone, but of deliberate human synthesis.

That led the team to suggest that particularly complex artificial molecules might serve as “molecular technosignatures”—chemical evidence oftechnology, rather than life itself.

For now, however, the method remains far from being a ready-to-fly alien-life detector. Instrument settings strongly affected performance. In simulated experiments, changing impact energies could nearly double prediction error, while one comparison increased it by more than 2.3 times. Molecules that break apart unusually easily can also cause the model to overestimate assembly, while molecules that resist fragmentation may be underestimated.

The researchers therefore describe the study as a proof of concept. Turning the approach into an effective tool for planetary investigation would require standardized databases, careful instrument calibration, and chromatographic separation to isolate individual molecules before analysis.

Future versions could also combine mass spectrometry with other techniques, including nuclear magnetic resonance or infrared spectroscopy. Infrared measurements, the researchers suggest, extend aspects of the approach to the remote study of exoplanet atmospheres.

Yet the wider goal is significant. Instead of deciding beforehand what alien biology should produce, scientists could search for evidence that a process has persistently directed chemistry toward highly specific, complex outcomes.

“In this work, we presented a proof-of-concept demonstrating the potential use of ML to predict the MA molecular complexity as a biosignature from in situ mass spectrometers during future space missions,” the researchers write. “Although the application of ML in life detection remains a nascent endeavor, it holds significant promise.”

The recent study, “Molecular Assembly as a Universal Biosignature Measurable by Mass Spectrometry,” appeared in Proceedings of the National Academy of Sciences.

Tim McMillan is a retired law enforcement executive, investigative reporter and co-founder of The Debrief. His writing typically focuses on defense, national security, the Intelligence Community and topics related to psychology. You can follow Tim on Twitter: @LtTimMcMillan.  Tim can be reached by email: tim@thedebrief.org or through encrypted email: LtTimMcMillan@protonmail.com 

A tiny microscope weighing just five grams could give neuroscientists an unprecedented view of how individual brain cells control behavior while animals move naturally. Researchers at the University of Colorado Anschutz Medical Campus and the University of Colorado Boulder have developed a head-mounted device that can both record neural activity and selectively activate individual neurons during movement. The technology could help scientists investigate how brain circuits generate behavior and how those circuits become disrupted in conditions such as Alzheimer’s disease, Parkinson’s disease and epilepsy.

The device, called Opto2P-FCM, combines two powerful techniques that have traditionally been difficult to use together in a freely moving animal. Two-photon microscopy provides detailed images of living brain tissue at cellular resolution, while optogenetics uses light-sensitive proteins to control selected neurons. By integrating both capabilities into a miniature instrument, the researchers can observe specific cells as they become active and then stimulate chosen neurons to test how they influence the surrounding network.

“These miniature microscopes have been increasingly used to study the neural basis of behavior in freely moving animals,” said Emily Gibson, PhD, associate professor of biomedical engineering at CU Anschutz and the study’s senior author. “This new microscope allows deeper penetration and higher resolution images in the brain. Importantly, it can also excite select neurons using optogenetics.”

For decades, conventional two-photon microscopes have been regarded as the gold standard for imaging neurons in living brains. These systems use infrared laser light to produce high-resolution images deep within tissue, allowing researchers to track calcium signals and other indicators of neural activity. However, conventional microscopes are generally large, stationary instruments. Animals must often remain immobilized beneath the objective, making it difficult to study brain activity during walking, exploration, social interaction or other natural behaviors.

Miniature microscopes have helped overcome that limitation by moving the imaging system onto the animal’s head. Yet reducing the size and weight of the equipment has often required compromises. Some compact systems provide lower-resolution images, while others can record neural activity but cannot stimulate precisely selected cells. The Colorado team designed Opto2P-FCM to address both problems through a dual-path optical architecture that separates imaging from stimulation.

In the imaging pathway, infrared light is directed through the brain to generate two-photon signals from fluorescent markers inside neurons. These markers can report changes in calcium concentration, which occur when neurons fire. The resulting signals are collected to create detailed images of neural activity at the level of individual cells. A separate optical pathway delivers patterned light for optogenetic stimulation, allowing researchers to activate specific neurons without disrupting the imaging beam or overwhelming the detectors.

This separation is technically important because imaging and stimulation place different demands on an optical system. High-resolution imaging requires precise focusing, efficient light collection and minimal optical distortion. Optogenetic activation, by contrast, requires intense, accurately positioned light capable of reaching the light-sensitive proteins expressed by selected neurons. Designing both functions into a device small enough to be carried by an animal required careful control of the alignment, geometry and weight of every optical component.

“This microscope is a game changer,” said Juliet Gopinath, PhD, professor of electrical, computer and energy engineering and physics at CU Boulder and a study co-author. “Being able to demonstrate a state-of-the-art instrument for both read-out and photo-stimulation of neurons is amazing.” The research team included engineers, physicists and neuroscientists who worked together to solve problems involving miniature lenses, mechanical tolerances, optical alignment and three-dimensional fabrication.

The device weighs approximately five grams, but its small size made construction unusually demanding. Mo Zohrabi, PhD, a senior research scientist at CU Boulder and co-first author, said the team initially thought such an instrument might be impossible to build because of the tight tolerances required by the optics and available 3D-printing technology. Gregory Futia, PhD, a senior research associate at CU Anschutz and co-first author, said that every component had to be designed and accounted for with extreme precision. Even slight misalignment could reduce image quality or prevent the stimulation beam from reaching the intended cells.

The researchers believe the platform could transform experiments designed to distinguish correlation from causation in the brain. Recording a neuron’s activity can show that it is associated with a behavior, but it does not prove that the cell helps produce that behavior. With optogenetic stimulation, scientists can activate a selected population and observe whether movement, decision-making or other actions change. At the same time, the microscope can record how neighboring neurons respond, providing a view of how local circuits coordinate activity.

The current Opto2P-FCM system is a prototype, and the researchers are working on future versions that could be smaller, lighter and faster. They also hope to expand the field of view so that more neurons can be monitored at once. Such improvements could make it possible to track larger networks over longer periods while animals engage in increasingly complex behaviors. By combining cellular-resolution observation with targeted manipulation in a freely moving subject, the technology may help reveal how healthy neural circuits operate and how neurological disease changes the communication between them.

Subject of Research: Miniature two-photon microscopy, optogenetics, neural circuits and brain activity during natural movement

Article Title: New Miniature Microscope Allows Scientists to Watch and Control Brain Cells During Natural Movement

News Publication Date: Aug. 11, 2026

Web References: University of Colorado Anschutz Medical Campus; University of Colorado Boulder; Optica study

References: Optica, “New Miniature Microscope Allows Scientists to Watch and Control Brain Cells During Natural Movement,” article publication date: Aug. 7, 2026

Keywords

Miniature microscope, two-photon microscopy, optogenetics, neuroscience, brain imaging, neural circuits, freely moving animals, biomedical engineering, applied physics, neurological disease, brain research