For the first time, scientists have used ultrafast X-ray flashes to take a direct image of a single electron as it moved during a chemical reaction.

In the new study, published Aug. 20 in the journal Physical Review Letters, the researchers accomplished this incredible feat by imaging how a valence electron — an electron in the outer shell of an atom — moved when an ammonia molecule broke apart.

For decades, scientists have used ultrafast X-ray scattering to image atoms and their chemical reactions. The scattering uses supershort bursts of X-rays to freeze tiny, fast-moving molecules in action. X-rays have the perfect wavelength range for capturing details at the atomic scale, which is why they're ideal for imaging molecules.

However, X-rays interact strongly only with core electrons near the atom’s nucleus. Valence electrons — the outermost electrons in an atom and the ones actually responsible for the chemical reactions — were hidden.

"We wanted to take pictures of the actual electrons that are driving that motion," Ian Gabalski, a physics doctoral student and lead author of the study, told Live Science.

If scientists can understand how valence electrons move during chemical reactions, it could help them design better drugs, cleaner chemical processes, and more efficient materials, Gabalski said.

To get started, the team needed to find the right molecule. It turned out to be ammonia.

"Ammonia is kind of special," Gabalski said. "Because it has mostly light atoms, there aren't a lot of core electrons to drown out the signal from the outer ones. So we had a shot at actually seeing that valence electron."

The experiment was conducted at the SLAC National Accelerator Laboratory's Linac Coherent Light Source, a facility that produces intense, short X-ray pulses. First, the team gave the ammonia molecule a tiny jolt of ultraviolet light, which made one of the electrons "jump" to a higher energy level. Electrons in molecules usually stay in low-energy states, and if they are pushed to a higher one, it triggers a chemical reaction. Then, with the X-ray beam, the researchers recorded how the electron's "cloud" shifted as the molecule began to break apart.

Related: The shape of light: Scientists reveal image of an individual photon for 1st time ever

In quantum physics, electrons aren't seen as tiny balls orbiting the nucleus. Instead, they exist as probability clouds, "where higher density means you're more likely to see the electron," Gabalski explained. These clouds are also known as orbitals, and each one has a distinct shape depending on the energy and position of the electron.

To map this electron cloud, the team ran quantum mechanical simulations to calculate the molecule's electronic structure. "So now this program that we use for these kinds of calculations goes and it figures out where the electrons are filling up those orbitals around the molecule," Gabalski said.

The X-rays themselves act like waves, and when they pass through the electron's probability cloud, they scatter in different directions. "But then those X-rays can go and interfere with each other," Gabalski said. By measuring this interference pattern, the team reconstructed an image of the electron's orbital and saw how the electron moved during the reaction.

They compared the results to two theoretical models: one that included valence electron motion, and one that didn't. The data matched the first model, confirming that they had captured the electron's rearrangement in action.

The researchers hope to adapt the system for use in more complex, 3D environments that better mimic real tissues. That would move it closer to applications in regenerative medicine, such as growing or repairing tissue on demand.

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

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

The first stars in the universe

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

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

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

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

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

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

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

Far more important than previously assumed

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

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

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

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

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

When lightning flashes above, plants on the ground may respond in kind. 

Scientists have long been aware that plants and trees can emit small, visible electric discharges from the tips of their leaves when the plants are trapped beneath the electrical fields generated by thunderstorms high overhead. These discharges, known as coronas, are sometimes visible as faint, blue sparks that glow around charged objects. 

Now, new research suggests those plant-based sparks may be altering the surrounding air quality in ways never recognized before. But whether the impacts of these minishocks in the atmosphere are positive or negative remains unclear.

In the study, published Aug. 9 in the Journal of Geophysical Research: Atmospheres, researchers recreated the electrical fields from thunderstorms in a laboratory and analyzed the coronas given off by eight plant species under a range of conditions. The results showed that all of the coronas created a high abundance of radicals — chemicals containing unpaired electrons that are highly reactive with other compounds — which can significantly alter the surrounding air quality. 

"While little is known about how widespread these discharges are, we estimate that coronas generated on trees under thunderstorms could have substantial impacts on the surrounding air," lead study author Jena Jenkins, an atmospheric scientist at Penn State University, said in a statement. 

Related: 'Gigantic jet' that shot into space may be the most powerful lightning bolt ever detected 

The two radicals given off by the plant coronas are hydroxyl (OH) and hydroperoxyl (HO2), both of which are negatively charged and are known to oxidize, or steal electrons from, a number of different chemical compounds, thereby transforming them into other molecules. The researchers were particularly interested in the concentrations of hydroxyl radicals because they have a greater impact on air quality.

"The hydroxyl radical contributes to the total atmospheric oxidation of many atmospheric pollutants," study co-author William Brune, a meteorologist at Penn State University, said in the statement. 

For example, if a hydroxyl radical reacts with greenhouse gases, such as methane, then it can remove the damaging molecules from the atmosphere and help combat climate change, Brune said. But if the same radical reacts with oxygen, it can create ozone, which, despite playing an important role in the upper atmosphere, is toxic to humans. The radicals can also create aerosol particles that harm air quality, he added. 

This is not the first time that researchers have shown the link between thunderstorms and hydroxyl radicals. 

In 2021, a research team led by Brune found that lightning was a major progenitor of hydroxyl radicals in the atmosphere. In their paper, published in the journal Science, the team theorized that thunderstorms could be directly responsible for up to one-sixth of the hydroxyl radicals in the atmosphere.

In September, another team led by Brune released a follow-up study, published in the journal Earth, Atmospheric and Planetary Sciences, that showed coronas produced by metallic objects such as telephone poles and transmission towers produce a slightly higher level of hydroxyl radicals than plant coronas. However, the levels of radicals produced by plant and artificial coronas are both significantly less than those produced directly from lightning. 

"Even though the charge generated by the [plant] corona was weaker than the sparks and lightning we looked at before, we still saw extreme amounts of this hydroxy radical being made," Jenkins said. 

Considering the vast numbers of trees that are present in lightning-prone areas, plant-produced coronas may represent a majorly understudied source of radicals with a highly unpredictable effect on air quality, she added.

"There are about two trillion trees in areas where thunderstorms are most likely to occur globally and there are 1,800 thunderstorms going on at any given time," Jenkins said.

As a result, researchers want to continue studying these coronas in greater detail to fully understand the effect they have on localized air quality and on a wider global scale. 

"The hydroxyl radical is the atmosphere's most important cleanser," Jenkins said. "So having a better accounting of where this stuff is being made can give us a more complete understanding of what's happening in the atmosphere."

Other studies suggest that thunderstorms may become more frequent and powerful due to the effects of human-caused climate change, so understanding the effects of thunderstorms on air quality is vital, she added. 

During the experiments, the team made another discovery that could help accelerate this field of research: The leafy discharges gave off sharp spikes of ultraviolet radiation. This could allow the team to indirectly study where coronas are occurring in the field and measure their effects on nearby air quality. 

Synthetic compounds known as "forever chemicals" because they never break down in the environment can actually be destroyed — by beheading. 

Scientists discovered a simple destruction technique that works on 10 types of these chemicals, known as per- and polyfluoroalkyl substances (PFAS). Researchers hope that the method will expose weaknesses in even more PFAS-class substances, leading to paths for removing these chemicals from drinking water easily and cheaply.

The researchers published their findings in the journal Science on Aug. 18.

Forever chemicals, everywhere

PFAS compounds are found in a huge variety of products: from food packaging to shampoos to non-stick cookware and electronics. The chemicals consist of a head, often containing charged oxygen molecules, and a tail of carbon and fluorine atoms, which are joined by a bond that nothing in nature can break. As a result, PFAS compounds persist in low levels in soil, air, water and even human blood samples, according to the Environmental Protection Agency (EPA). Researchers are still working to understand the health effects of PFAS exposure, but the chemicals may interfere with the body's hormones, raise cholesterol levels, affect fertility and increase the risk of certain cancers, according to the EPA.

Related: What are PFAS?

PFAS chemicals pass through water treatment plants unaltered, and current clean-up methods are prohibitively expensive, Tasha Stoiber, an environmental chemist at the Environmental Working Group, a nonprofit organization that tracks environmental contamination, told Science. Stoiber was not involved in the new study. Filtering them and sticking PFAS chemicals in landfills just raises the risk that they'll leach out later, and other clean-up methods often produce harmful byproducts that require complicated steps to break down. 

Chemists at the University of California, Los Angeles and Northwestern University took a simpler approach — targeting the molecules' heads with a chemical guillotine. Charged oxygen molecules are reactive with other chemicals, so Northwestern University chemist William Dichtel and doctoral student Brittany Trang decided to go after this weak spot. 

The researchers used a common solvent, dimethyl sulfoxide (DMSO), because previous EPA research had hinted that PFAS slowly degrades when exposed to DMSO. Testing different recipes at different temperatures, Dichtel and his colleagues discovered that the fastest way to take PFAS down was to heat the "forever chemical" to boiling along with DMSO and lye, or sodium hydroxide — a common chemical found in many types of soap. 

"That triggered all these reactions, and it started spitting out fluorine atoms from these compounds to form fluoride, which is the safest form of fluorine," Dichtel said in a statement. "Although carbon-fluorine bonds are super-strong, that charged head group is the Achilles' heel."

Cleaning up the water 

Fluoride is safe for humans in small amounts and is often added to drinking water to help prevent tooth decay. The other byproducts of the reaction are carbon dioxide and formic acid, which is a chemical used for defense by some species of ants and is low in toxicity — the acid is even sometimes added to animal feed to reduce bacterial growth without harming livestock.  

Because the new method doesn’t produce harmful chemicals, it could be used to clean drinking water, Trang told Science. PFAS compounds could be filtered away and then beheaded separately, leaving behind only harmless chemicals to dispose of. 

The method can clean up any PFAS compounds containing oxygen-rich carboxylic acid heads, the researchers found. That doesn't cover all classes of PFAS, but Dichtel and his colleagues are hopeful that their method might inspire new means of attacking more resistant PFAS chemicals. 

Originally published on Live Science.

Scientists recreated the unique chemical conditions found on Titan, Saturn's largest moon, in tiny glass cylinders here on Earth, and the experiment revealed previously unknown features of the moon's mineral makeup.

Titan is the second-largest moon in the solar system, behind Jupiter's Ganymede, and sports a dense atmosphere of mostly nitrogen with a dash of methane, according to Space.com. This yellowish haze hovers around minus 290 degrees Fahrenheit (minus 180 degrees Celsius). Below the atmosphere, lakes, seas and rivers of liquid methane and ethane cover Titan's icy crust, particularly near the poles. And similar to liquid water on Earth, these natural gases take part in a cycle in which they evaporate, form clouds and then rain down on the moon's surface.

Titan's dense atmosphere, surface liquid and seasonal weather cycles make the frigid moon somewhat similar to Earth, and like our planet, the moon is known to have organic molecules that contain carbon, hydrogen and oxygen, according to NASA. Because of this organic chemistry taking place on Titan, scientists think the moon could serve as a massive laboratory to study chemical reactions that occurred on Earth before the emergence of life on the planet, Space.com previously reported. 

Related: Moon birth and methane weather: Cassini's 7 oddest Saturn finds 

But only one spacecraft, Cassini, has observed Saturn and its moons in detail, making it tough to do Earthbound research on the wacky chemistry found on Titan. So recently, a team of scientists set out to simulate Titan in a test tube. 

The team first placed liquid water in small glass cylinders and cranked down the temperature to Titan-like conditions, the researchers said in a statement. This water froze to mimic Titan's icy crust. The team then introduced ethane to the tube, which became liquid like the lakes on Titan's surface. Finally, they added nitrogen to stand in for Titan's atmosphere and then varied the temperature of the tube ever so slightly, to simulate the variations in temperature on Titan's surface and in different layers of its atmosphere.   

In their recent study, presented Thursday (Aug. 26) at the fall meeting of the American Chemical Society, the team then added two compounds, called acetonitrile (ACN) and propionitrile (PCN). Data from the Cassini mission suggest that these compounds are abundant on Titan, principal investigator Tomče Runčevski, an assistant professor in the Department of Chemistry at Southern Methodist University in Dallas, told Live Science.

Most previous studies examined these two compounds separately, in their pure forms, but Runčevski's team wanted to see what would occur when the compounds mixed and mingled, as they might on Titan. As opposed to working with each compound separately, "if you mix them together ... there might be a completely different outcome in structure, so how the molecules will organize, and how the molecules will crystallize," or phase into a solid form, Runčevski said. 

And the team found that, when both present in Titan-like conditions, ACN and PCN behave quite differently than either compound in isolation. Namely, the temperatures at which the compounds melted or crystallized shifted drastically, on the order of tens of kelvins (hundreds of degrees Fahrenheit or Celsius).

Related: 6 most likely places for alien life in the solar system

These melting and crystallization points would be relevant in Titan's hazy yellow atmosphere. The various layers of the atmosphere differ in temperature depending on their altitude above the moon's surface, so to understand how chemicals behave throughout the haze, the new study suggests that these temperature variations need to be taken into account, Runčevski said.

In addition, the team found that, when ACN and PCN crystallize, they adopt different crystal structures depending on whether they're alone or in the presence of the other compound. Crystals form when the individual molecules within a compound snap into a highly organized structure. While the building blocks of that structure — the molecules — remain the same, depending on factors such as temperature, they can end up snapping together in slightly different configurations, Runčevski said.

These variations in crystal structure are known as "polymorphs," and when on their own, ACN and PCN adopt one polymorph at high temperatures and another at low temperatures. But "what we notice is that if we have a mixture, the stability of the high-temperature and low-temperature [polymorphs] can be, in a way, switched," Runčevski said. 

These fine details of when and how the compounds achieve a stabilized structure "can really change our understanding of what kind of minerals we might encounter on Titan," in terms of what polymorphs they likely adopt on the moon, he said. This in turn can shape what chemical reactions take place between these and other compounds on Titan.

The new study is limited in that it doesn't account for all of the chemicals present on Titan, and so can capture only a simplified picture of what actually happens on the moon, Runčevski said. 

"It's important for us as scientists on Earth ... to create these models with increasing complexity, and one day to reach models that are really significant and can really help us further understand the surface of Titan," he said.

NASA's Dragonfly mission, set to launch in 2026 and arrive at Saturn in 2034, may provide more on-the-ground information about the mineral makeup of Titan. However, Runčevski suspects that the crystals his team has observed likely form around the edge of Titan's lakes, cropping up as the liquid ethane in the lakes evaporates and leaves those compounds behind on the shoreline. At this point, it's unclear whether the Dragonfly mission might focus on this specific aspect of the Titanian environment, but "nonetheless, [the mission] is super exciting, and we will learn so much more about Titan," he said.

Originally published on Live Science.

A crumbling hunk of rock found in a field in England is a rare meteorite from the earliest days of the solar system, dating back about 4.6 billion years. 

The meteorite was found in Gloucestershire in March by Derek Robson, a resident of Loughborough, England, and the director of astrochemistry at the East Anglian Astrophysical Research Organisation (EAARO). The meteorite was sitting in the imprint of a horseshoe left behind in a field, according to Loughborough University.

The space rock is a carbonaceous chondrite, a rare category that makes up only 4% to 5% of meteorites that are found on Earth. These meteorites hail from the asteroid belt between Mars and Jupiter and formed early in the history of the solar system. Intriguingly, they often contain organic, or carbon-bearing, compounds, including the amino acids that make up the basic building blocks of life. This raises questions about whether these meteorites hold clues to how living things first emerged in the solar system. 

Related: The 7 strangest asteroids: Weird space rocks in our solar system 

Unlike other space debris, this chunk of rock didn't endure the violent collisions and intense heat involved in the creation of the solar system's planets and moons.

Rather, the meteorite has "been sitting out there, past Mars, untouched, since before any of the planets were created," Shaun Fowler, a microscopist at Loughborough University, said in a statement, "meaning we have the rare opportunity to examine a piece of our primordial past."

The rock is small, charcoal-colored and fragile, sort of like a chunk of crumbling concrete. The meteorite is mostly made of minerals such as olivine and phyllosilicates, Fowler said, as well as round grains called chondrules, which were partially molten beads incorporated into the asteroid when it first formed.

"But the composition is different to anything you would find here on Earth and potentially unlike any other meteorites we've found — possibly containing some previously unknown chemistry or physical structure never before seen in other recorded meteorite samples," Fowler said.

Researchers at Loughborough University and EAARO are using electron microscopy to study the surface of the meteorite down to the nanometer (a billionth of a meter), as well as techniques called vibrational spectroscopy and X-ray diffraction, which allow them to delve into the chemical structure of the minerals in the meteorite. If the team can confirm the presence of amino acids in the sample, the findings might reveal new information about how the early geochemistry of the solar system set the stage for life. The examination of the meteorite is still in the initial stages. 

"At this stage, we have learned a good deal about it, but we've barely scratched the surface," Sandie Dann, a chemist at Loughborough University, said in the statement.

Originally published on Live Science.

This Friday (July 9), we will explore the chemistry of milk plastic, (or casein plastic) in our new kids video series: Summer School with Live Science.

In this week's installment, Live Science producer, Diana Whitcroft, will demonstrate how to make biodegradable plastic, called casein plastic, with milk and vinegar. She will make jewelry in this tutorial, but families are urged to make anything they’d like, be it ornaments, keepsakes, decorations ... anything!

Every Friday at 3 p.m. EDT (12 p.m. PDT), Diana will host Summer School With Live Science, which you can find live on Live Science's Facebook, YouTube and Twitter pages. Every week, the series will explore a different field of STEM (science, technology, engineering and mathematics) through simple hands-on experiments that you and your child can follow along with at home. 

Disclaimer: It is strongly advised that all science experiments, recipes and methods be attempted only under adult supervision. Adults are required to handle or assist with any potentially harmful utensils and ingredients. Always wash hands thoroughly after trying any experiment. Avoid touching your face and eyes when performing any experiments, and if possible, wear glasses or safety goggles. Do not ingest any of the ingredients during or after performing this experiment.

Milk plastic: Materials

Age Range: 4-10 years

• 1 cup whole milk
• 4 Tbsp. white vinegar
• Small saucepan
• 1 large mug
• 3-4 paper towels
• Food coloring, beads, glitter (optional)
• Cookie cutter (optional)

Step One: Choose Your Design

In Diana’s tutorial, she will show you how to make a ring out of casein. But you and your kids are encouraged to make any decorative item you desire, such as ornaments or keepsakes. Feel free to double the recipe and make larger objects. Embellish with beads or glitter, too!

Step Two: Make Casein Plastic

Warm up the milk in a small saucepan. Do not bring to a boil. While you're waiting for your milk to heat, add the vinegar to your large mug. If doubling the recipe, think about upgrading from a mug to a mixing bowl. Once the milk is only just starting to bubble, remove it from the heat and add it to your mug of vinegar. Right away, you'll notice curdling, but allow your mug to cool for 3-5 minutes before handling.

Step Three: Mold Your Object

With a spoon, do your best to separate the milk solids from the liquid and transfer to a paper towel. Knead for about a minute before adding your color and decorations. Continue to knead, gently pressing what is now a dough against your paper towel so as to let out as much liquid as possible. When your dough is dry enough, it should resemble a play-dough type consistency. Shape it however you want. 

Document this experience and send images to us either on social media or to community@livescience.com. We'd love to see your results so that we can feature them in our photo gallery!

What's so important about casein plastic?

Milk contains the molecules of a protein called casein. When the milk is warmed, and acid added to it, (vinegar in this case), the casein molecules unfold and reorganize into a chain of monomers known as a polymer. If the process of cheesemaking comes to mind as you attempt this experiment, you’d be onto something, as the processes are very similar! In fact, the word casein is Latin for "cheese" and the wet substance you'll use to mold into jewelry looks a little bit like cottage cheese. This polymer will harden within 24 hours; and although it's not as durable as most industrial plastics, casein is biodegradable and quite versatile.

Casein plastic was first introduced in the early 1900s and is still used today to make things like buttons, buckles, jewelry, fountain pens, beauty accessories and so much more. To further reinforce the structural integrity of this plastic, many manufacturers have added the step of soaking it in formalin, which is a 5% solution of formaldehyde in water. Casein plastic was used to make jewelry for Queen Mary I of England, who ruled England from A.D. 1553-1558.

Originally published on Live Science.

This Friday (June 18), we will explore the critical field of engineering in our new kids video series: Summer School with Live Science. 

In this week's installment, Live Science producer, Diana Whitcroft, will demonstrate how to construct a capsule that will securely land a raw egg without breaking it. She will provide her own design, intended to protect this delicate payload from a mere 4- to 5-foot drop. Watch along with your kids and get inspired to engineer an even better contraption meant to land from even greater heights!

Every Friday at 3 p.m. EDT (12 p.m. PDT), Diana will host Summer School With Live Science, which you can find live on Live Science's Facebook, YouTube and Twitter pages. Every week, the series will explore a different field of STEM (science, technology, engineering and mathematics) through simple hands-on experiments that you and your child can follow along with at home. 

Disclaimer: It is strongly advised that all science experiments, recipes and methods be attempted only under adult supervision. Adults are required to handle or assist with any potentially harmful utensils and ingredients. Always wash hands thoroughly after trying any experiment. Avoid touching your face and eyes when performing any experiments, and if possible, wear glasses or safety goggles. Do not ingest any of the ingredients during or after performing this experiment.

Egg Drop Challenge: Objective

Age Range: 6+ years

Construct a capsule or package to secure and successfully land a raw egg, unfractured from a fall to the ground.

Step One: Make Blueprints

If you plan to experiment with a landing capsule of your own design, (and you really should), sketch it out! First, determine the height from which you plan to drop your egg. Then, assess the kind of materials available to you and go from there. Draft as many designs as you'd like and select one to build. 

Keep in mind that it is completely possible your egg might not survive the fall. Don't be discouraged by this. Simply work with another of your ideas until you land a happy, fracture-free huevo. After all, we can learn just as much through our failures as we do our successes.

Most students prefer the idea of encasing their eggs within enclosed capsules like shoe boxes or tupperware. Many, looking for a loftier egg drop, lean toward creating the force of drag by the use of parachutes. In Diana's demonstration, she set out to use the fewest and simplest materials possible, by creating a paper straw barrier to protect the egg from its 4-foot journey to the floor.

Step Two: Build Your Landing Capsule

How you go about constructing your piece is entirely up to you and is dependent upon the complexity of your contraption. Log your methods, so that you and your child can determine why your drop was successful (or unsuccessful). In the case of our paper straw capsule, Diana found that by using paper straws instead of plastic straws, and by first taping the straws against the egg with clear tape and then reinforcing the connecting parts with electrical tape, she was able to engineer a more encompassing and structurally sound device.

Step Three: Drop It Like It's Hot

After you've designed and engineered your device to withstand the desired drop height, be sure to use measuring tape to figure out roughly how high to hold your egg before releasing it. This experiment is better performed outside. To avoid wasting an egg or making a mess, set up a drop cloth or a tray to catch any spatter. Wear protective eye gear and either a smock or apron. To better record your results, try setting up a camera to record the drop. 

Step Four: Log Your Results

Whether this experiment was a success or a yolky failure, write it down. Make this a summer-long project if you'd like by constructing a series of contraptions to see how high you can successfully land an egg, or even go head-to-head with family members to see who can make the best landing capsule. The more creative you get with your designs, the more you'll learn.

Document this experience and send images to us either on social media or to community@livescience.com. We'd love to see your results so that we can feature them in our photo gallery!

Why Perform An Egg Drop?

Have you ever watched NASA's video illustration of the landing of Mars rovers, Spirit and Opportunity? Compare those landings to that of Perseverance. Have you ever wondered how we deliver aid and general goods to devastated or hard-to-access areas of the world? This activity aims to give kids an idea of how we achieve these goals.

By engaging in this highly versatile activity, kids are effectively simulating these real-life applications and in doing so, learning the importance of design iteration through trial and error. It is an exceptional way to introduce engineering design to young minds.

General physics and materials science are also highlighted in this experiment. Whether you and the family choose to drop a straw capsule from a 4-foot height or send your egg off with a parachute from a two-story landing, students are challenged to engineer a device that would decrease the amount of impulse of the egg that gets transferred from potential energy into kinetic energy on the egg's shell. When objects collide, (in this case, the egg with the ground), their energy and momentum are changed or transferred. These processes are mediated by one or multiple forces — the force that gets the egg moving and transfers its potential to kinetic energy is … gravity. Then upon contact with the ground or surface, the kinetic energy must be transferred somewhere and that’s the point of impact.  If the forces acting on the egg are too strong, that can cause the shell to fracture.  Thus the development of these different engineered structures to divert this energy from the sensitive payload or reduce the kinetic energy altogether (i.e. parachutes).

Originally published on Live Science.

This Friday (June 11), we will explore the eruptive realm of acid-base chemistry in our new kids video series: Summer School with Live Science.

In this week's installment, Live Science producer, Diana Whitcroft, will demonstrate the reaction between sodium bicarbonate (baking soda) and citric acid (lemon juice, in this case). This fun experiment is a great way to introduce young minds to the world of chemistry.  

Every Friday at 3 p.m. EDT (12 p.m. PDT), Diana will host Summer School With Live Science, which you can find live on Live Science's Facebook, YouTube and Twitter pages. Every week, the series will explore a different field of STEM (science, technology, engineering and mathematics) through simple hands-on experiments that you and your child can follow along with at home. 

Disclaimer: It is strongly advised that all science experiments, recipes and methods be attempted only under adult supervision. Adults are required to handle or assist with any potentially harmful utensils and ingredients. Always wash hands thoroughly after trying any experiment. Avoid touching your face and eyes when performing any experiments, and if possible, wear glasses or safety goggles. Do not ingest any of the ingredients during or after performing this experiment.

Lemon Volcanoes: Materials

Age Range: 5-10 years

• Tray or Cookie sheet
• Paper (optional)
• Lemons
• Cutting knife
• Popsicle stick
• Spoon
• Food coloring
• Dish soap
• Baking soda

Step One: Prep Your Work Station

Place your tray on a surface that has been well-covered with newspaper. Optional: Place paper on your tray; this is where you'll be erupting your volcanoes and if the paper catches the colorful liquid, you can turn it into artwork.

Step Two: Ready Your Lemon

Have an adult cut a small portion of the bottom off the lemon so that it stands upright. Then, cut off the top of the lemon, exposing the inside pulp and juice. Using your popsicle stick, mash the inside of the lemon so that the "meat" or fleshy interior is completely pushed down and the lemon is filled with mostly juice. Be careful not to puncture the lemon rind.

Step Three: Choose a Color

Squeeze a few drops of your food coloring into your lemon. The amount of food coloring you use is dependent on the size of your lemon and how much juice it yields. Small lemons with very little juice: 2-3 drops. Large lemons with a lot of juice: 4-5 drops. Try different color combinations and see what comes out! You can even erupt two or three lemons at a time for a tie-dye effect on your paper. Then, give a good squirt of dish soap into your lemon.

Step Four: Erupt Your Volcano

Using your spoon, sprinkle into your lemon a few teaspoons of baking soda. Again, depending on the size of your lemon, you'll need to adjust the amount. Small lemons with little juice: 2-3 tsp. Large lemons with a lot of juice: 4-5 tsp. Right away, you should see bubbling and fizzing, but it is likely that you'll need to use your spoon or popsicle stick to better mix the juice and baking soda. Your lemons will now start to overflow with colorful bubbles. 

Document this experience and send images to us either on social media or to community@livescience.com. We'd love to see your results so that we can feature them in our photo gallery!

Step Five: Paint!

Now that the explosive fun is over, it's time to get creative. Using your hands, a paint brush or even sponges, create your own works of art from the colorful liquid left behind on the paper. Simply remove your lemons and go to town! Be sure not to over-paint your paper. Combining your colors too much will turn your colorful canvas into a big grey blob, (unless of course you really like the color grey!).

The science behind lemon volcanoes

When the baking soda was added to the lemon juice, it bubbled and foamed. That is because when sodium bicarbonate (baking soda) and citric acid (lemon juice) are combined, they react by forming carbon dioxide gas as well as a chemical compound called sodium citrate. It's the carbon dioxide that causes all those bubbles (remember CO2 is the gas that makes your can of soda or seltzer water fizz). We know that carbon dioxide is the primary gas we exhale when breathing. It occurs naturally in Earth's atmosphere and is emitted by natural sources like geysers (and volcanoes) as well as industrial processes. 

Sodium citrate is the sodium derived from citric acid. It is used in a variety of everyday items. It is used as an anticoagulant for blood products, as well as an acidity regulator in food.

Scientists have discovered a new type of uranium that is the lightest ever known. The discovery could reveal more about a weird alpha particle that gets ejected from certain radioactive elements as they decay.

The newfound uranium, called uranium-214, is an isotope, or a variant of the element, with 30 more neutrons than protons, one fewer neutron than the next-lightest known uranium isotope. Because neutrons have mass, uranium-214 is much lighter than more common uranium isotopes, including uranium-235, which is used in nuclear reactors and has 51 extra neutrons. 

This newfound isotope isn't just lighter than others, but it also showed unique behaviors during its decay. As such, the new findings will help scientists better understand a radioactive decay process known as alpha decay, in which an atomic nucleus loses a group of two protons and two neutrons — collectively called an alpha particle.

Though scientists know that alpha decay results in the ejection of this alpha particle, after a century of study, they still don't know the exact details of how the alpha particle is formed before it gets ejected. 

Related: Elementary, my dear: 8 elements you never heard of

The researchers created the new uranium isotope at the Heavy Ion Research Facility in Lanzhou, China. There, they shone a beam of argon at a target made of tungsten inside a machine called a gas-filled recoil separator — in this case the Spectrometer for Heavy Atoms and Nuclear Structure, or SHANS. By shining a laser at the tungsten, the researchers effectively added protons and neutrons to the material to create uranium.

The new uranium-214 isotope had a half-life of just half a millisecond, meaning that's the amount of time it takes for half of the radioactive sample to decay. The most common isotope of uranium — called uranium-238 — has a half-life of about 4.5 billion years, which is about the age of Earth. 

By carefully watching how the isotopes decayed, the scientists were able to study the strong nuclear force — one of the four fundamental forces that hold matter together — acting on the alpha particle parts — the neutrons and protons — on the surface of the uranium. They found that the proton and neutron in each alpha particle interacted much more strongly than in isotopes and other elements with similar numbers of protons and neutrons that have been previously studied. 

This is likely due to the specific number of neutrons inside the nucleus of uranium-214, the researchers said. The new isotope has 122 neutrons, nearing the "magic neutron number" of 126, which is especially stable due to the configuration of the neutrons in complete sets, or shells. With this configuration, it is easier for scientists to calculate the strong force interaction between the protons and neutrons. That makes these isotopes particularly interesting to scientists, since studying these interactions can reveal features related to nuclear structure and decay process, said study lead author Zhiyuan Zhang, physicist at the Chinese Academy of Sciences. 

The scientists suspect that this proton-neutron interaction could be even stronger heavier radioactive elements such as isotopes of plutonium and neptunium. These elements have a few more protons, and the configuration of their orbits suggests they could have even stronger interactions than the uranium isotopes. . The scientists would like to study other elemental isotopes near the magic neutron number; however, since such elements have even shorter half-lives, even more sensitive detectors and more powerful beams will be needed.

The new findings were published April 14 in the journal Physical Review Letters.

Originally published on Live Science.

Science savvy individuals deserve a gift that speaks to their nerdy side, honoring that calculated love with thoughtful mementos. For recipients with a chemistry habit, the best place to begin is with the elements, as in the periodic table of elements. This chemistry classic appears on everything from decor to clothing. So, there's something to fit every personality. We've collected the perfect periodic table gift options to excite the chemistry aficionado in your life.

Periodic presents

Finding a superb gift for the chemist in your life doesn't have to be difficult, just begin with the periodic table. This colorful element-filled chart is the star of many gift-worthy items, designed to inject a little science flavor into homes and even wardrobes. Our staff pick is the Sunlit updated periodic table of elements shower curtain because it splashes color and personality into any bathroom, transforming ordinary tub and shower combinations into scientific sanctuaries. This standard-size curtain is anything but, with a full-color print of the periodic table. It's shaped in easy polyester and cleans in the washing machine. This gift is a no-brainer.

For those just beginning their love for chemistry, the Uncle Goose periodic table blocks set is a great foundational choice. Future chemists will enjoy this set of 20 elemental cubes, configuring periodic charts as they stack and play. Each side of the 20 wooden cubes is printed with an element, so babies will be introduced to what chemistry has to offer.

Periodically you'll have to buy a gift for the science enthusiast in your life, specifically a chemist or chemist-to-be. (If you are that person, you just might have to treat yourself.) To find the perfect gift, you'll need to consider everything, right down to the molecule. The gift has to possess a balanced combination of elements: chemistry flair, humor, and function. To help you get just the right reaction, we've compiled the very best gifts for chemistry lovers.

Give the gift of science

Fans of all things chemistry deserve a sciencey gift from time to time. These gifts should be equal parts fun and functional, playfully commemorating the field they know and love. There are a variety of creative options out there, but the best place to start is with clothing. The Ann Arbor T-Shirt Co. never trust an atom t-shirt playfully picks on the smallest unit of matter with a cheeky atom print and humourous phrase. This comfortable tee is shaped in lightweight cotton and is available in both women's and men's sizing.

For chemists who are coffee enthusiasts, the American Scientific caffeine beaker mug offers a bit of laboratory fun to that time-honored morning beverage. This mug emulates scientific glassware, taking the shape of a standard beaker. It features a caffeine molecule graphic and milliliter measurements so coffee consumption can be tracked.  

Scientists may finally understand the mysterious transition behind a century-old chemistry experiment. The details of this transformation, in which adding electrons to a bright blue ammonia solution morphs it into a lustrous, metallic bronze, have long eluded scientists.

The new study reveals the subtle details of this change, and shows that this transformation is gradual, rather than sudden. "What we've done successfully is that we've pretty much understood how these solutions behave at a wide range of concentrations using a microjet technique," said study co-author Ryan McMullen, a doctoral student in chemistry at the University of Southern California. This technique, which involves shooting hair-thin streams of the solution through a vacuum, has not been used on the lustrous liquid before.

And the discovery could open up new types of reactions in organic chemistry in the future, McMullen told Live Science.

Related: 8 chemical elements you never heard of

What is a metal?

Metals are a diverse group. Some, like lithium, are light enough to float, while others, like lead or osmium are extremely dense. Some require incredibly high temperatures to melt, while others melt easily (Mercury, for example, melts at minus 38.3 degrees Celsius, or minus 37.9 degrees Fahrenheit). Ultimately, what metals have in common is their ability to conduct electricity at absolute zero, the point at which molecular movement from heat essentially halts. 

But how do some nonmetals transform into metals? In a new study, researchers answered that question by adding metals to liquid ammonia. 

First, the researchers condensed ammonia, which is a gas at room temperature, into a liquid by cooling it to negative 27.4 F (minus 33 C). They then added either sodium, lithium or potassium, which are all alkali metals. (Rather famously, these metals react explosively when submerged in water.) The experiments were done in collaboration with scientists from the Czech Academy of Sciences and the Fritz-Haber Institute of the Max Planck Society in Berlin, as well as researchers in Japan and France.

Related: The top 10 greatest explosions ever

The result was an expected reaction: The liquid ammonia pulled electrons from the metal. Those electrons then became trapped between the ammonia molecules, creating the so-called solvated electrons the researchers hoped to study. At low concentrations, the result was a blue, non-metallic liquid. As the solvated, or trapped, electrons piled up, though, the solution transitioned to shiny bronze.

The next challenge was to investigate how the solvated electrons behaved at different concentrations. This involved shooting a microjet of the solution — about the width of a human hair — through a beam of synchrotron X-rays, which are high-energy X-ray beams. The X-rays excited the solvated electrons, causing them to hop out of their liquid cage of ammonia molecules. The researchers could then measure how much energy it took to release the solvated electrons.

The researchers found that the greater the concentration of solvated electrons, the more the pattern of energy release matched what is seen in a metal. Here's what that means: If you graph the amount of energy required to free electrons from their liquid ammonia cage, metals typically have what's called a "Fermi edge," a very abrupt transition, McMullen said. At lower concentrations of solvated electrons, this energy-release graph looks more like a rounded hill. Only at higher electron concentrations did this Fermi edge emerge. The edge reflects how much energy electrons have at a given temperature, McMullen added.

"When you increase the concentration to the metallic range then you see, this wonderful pattern emerges that is very, very characteristic of a metal," McMullen said.

The results were interesting because they showed that the metal-like liquid created by combining alkali metals and ammonia actually is a metal on a fundamental physical level, he said.

"It is a genuine metal, it's not something that just looks like one," McMullen said.

Lower-concentration solvated electrons are used in a type of reaction called a Birch reaction, which adds electrons to molecular structures called aromatic rings. This kind of reaction was used in the manufacture of the first oral contraceptive pills in the 1950s, McMullen said. By understanding how solvated electrons work at high concentrations, researchers can potentially find new kinds of chemical reactions, he said. For example, they might excite the solvated electrons with beams of light to get them to behave in new ways.

"If you tickle the electrons a bit so that they're more energetically excited, you can start looking at some crazy reactions that would never otherwise happen," McMullen said.

The researchers reported their findings June 5 in the journal Science.

Originally published on Live Science.

An instructor and two children pour three cups of powder into a bin of red liquid. Suddenly — poof — a cloud of what looks like cotton candy explodes toward the ceiling.

This popular video on Twitter comes courtesy of the Malay-language account w, which shares science content. But what is going on in the video?

It's a rather dangerous version of a classic chemistry demonstration, according to Brian Hostetler, an educator at the Denver Museum of Nature & Science. The reaction is typically known as "elephant's toothpaste," due to its foamy appearance, and it's commonly used in chemistry classrooms to explain catalysts, Hostetler told Live Science. [Elementary, My Dear: 8 Little-Known Elements]

Easy but explosive

The reaction uses cheap, easy-to-access ingredients: hydrogen peroxide, dish soap, potassium iodide and food coloring. Hydrogen peroxide is key. It's made up of two hydrogen and two oxygen molecules. The bonds between these molecules naturally break, so over time, hydrogen peroxide slowly becomes water and oxygen gas. That reaction happens faster when exposed to light, Hostetler said, which is why hydrogen peroxide is sold in brown bottles.

Normally, the slow breakup (or decomposition, in chemistry terms) of hydrogen peroxide into water and oxygen is unnoticeable. But the elephant's toothpaste experiment speeds the process with a catalyst, a chemical compound that increases the rate of a given reaction. Potassium iodide — a salt of iodine, and the dietary supplement that's used to add iodine to table salt — provides that catalyst.

"In the presence of potassium iodide, hydrogen peroxide decomposes almost immediately," Hostetler said.

The setup is simple. Hydrogen peroxide is mixed with dish soap, and food coloring is often added for a dramatic effect (which explains the cotton-candy pink in the Twitter video). The potassium iodide is added, and the iodide ion that's part of that compound attracts the oxygen in the hydrogen peroxide, breaking the bonds and rapidly transforming the hydrogen peroxide into water and oxygen gas. The oxygen molecules then get trapped by the soap, forming bubbles, Hostetler said. In a step sometimes added to the elephant's toothpaste demonstration, a glowing splint — a strip of wood that is hot but not burning — that is inserted into the bubbles will catch alight, sparked by the pure oxygen.

Toning it down

Usually, Hostetler said, the elephant's toothpaste experiment creates an oozy concoction. So why did the Twitter version send bubbles flying toward the ceiling?

That particular reaction was due to the strength of the ingredients and the shape of the containers, Hostetler said. A fairly safe version of the elephant's toothpaste demonstration can be done at home with 3% hydrogen peroxide bought from the drugstore, with yeast as the catalyst (yeast contains the enzyme catalase, which also breaks down bonds in hydrogen peroxide). The combination will ooze and get a little warm as the reaction releases heat, but other than the need to take care not to touch the "toothpaste," as hydrogen peroxide can be irritating to skin and eyes, this DIY version is pretty safe.

The Twitter video probably shows the reaction with 30% hydrogen peroxide, or even stronger, Hostetler said. The demonstrator also uses potassium iodide that's in powder form instead of mixed into water. And he has it poured in three batches at once into a large container with a lot of surface area, so the reaction happens across a large amount of hydrogen peroxide all at one time.

That makes the scene in the video "super-duper dangerous," Hostetler said. Thirty percent or higher hydrogen peroxide can cause chemical burns on the skin, he said, and the reaction could heat the solution by hundreds of degrees. It's that heat and steam from the reaction that floats some of the foam skyward in the Twitter video.

The bottom line, Hostetler said, is not to try the jumbo version of the demonstration at home – but feel free to press "play" again on Twitter.

"It's a cool video," he said.

• Top 10 Greatest Explosions Ever
• Creative Genius: The World's Greatest Minds
• Cotton Candy Grapes: The science behind the sweet, carnival taste

Originally published on Live Science.

A new video on imgur shows two kids peeking out of a towering "bubble igloo" emerging from a bathtub. The column of bubbles reaches nearly to the ceiling, leaning slightly but not collapsing. "HOW???" asked one commenter, quite reasonably.

The original poster hasn't clarified precisely how they created the tower of bubbles, but you can probably do this at home … if you don't mind a mess. Surfactants — like soap — can easily form stable foams, said Donald Freese, a chemical engineer formerly at W.L. Gore. Foams are gas bubbles dispersed in liquid; no foam is perfectly stable, because bubbles eventually pop and the gas pockets merge, becoming larger and breaking down the overall foam structure. But when soap bubbles are small and there are few disturbances, foams can stick around and stand up against the forces of gravity for a long time, Freese said. [Check Out Amazing GIFs of Chemical Reactions]

"Bubbles do pop, but the rate that they pop can be very slow for some types of surfactants if they are concentrated enough and if the system is free of oily contaminants and any other disrupting forces, such as airflow," Freese said.

Stable foams are frequently used in firefighting, Freese said. They're also used as a drilling fluid in oil wells, according to oil field company Schlumberger; special foams are sometimes added to the well as the drill bit churns through the rock to lubricate it and facilitate the movement of dust and rock chips. Stable foams might even show up in the kitchen, as when beating egg whites into a stiff meringue.

A bubble tower like the one in the imgur video could be made with soapy water (from dish soap or bubble bath) and an air pump to introduce the gas, said Laurent Courbin, a CNRS researcher at the Physics Institute of Rennes at the University of Rennes 1 in France.

"New bubbles would then be created at the level of the liquid surface, while old ones on top of the foam get higher," Courbin told Live Science. A barrier like a shower curtain, pulled away before filming started, might have helped to keep the bubbles corralled.

As for how to churn up enough air to raise bubbles to the ceiling, bathtub jets might do the trick. For the tragically jetless, some companies sell bubble machines for bathtubs. These bubble makers are shaped like mats that go on the bottom of the tub or units that attach to the side, and they generally cost around $100 or less.

Original article on Live Science. 

A way-cool clip posted on Reddit shows the strange phenomenon of a "snap freeze."

Poster DifficultBoss shared a short video of a bottle of spring water rapidly freezing, from top to bottom, when shaken after being left outside in cold weather. The science behind this sudden transition has to do with how water changes from liquid to solid.

When it hits its freezing point of 32 degrees Fahrenheit (zero degrees Celsius), water normally starts to crystallize into ice. But this crystallization process requires a home base — a nucleus that water molecules can arrange themselves on. If the water is pure enough, crystallization won't happen, according to the University of Illinois Physics Van outreach program. Liquid water that is below freezing is said to be "supercooled." [Check Out Amazing GIFs of Chemical Reactions]

The snap freeze happens when this delicate, supercooled state is disturbed, as it is in the video when the poster gives the bottle a small shake. The shake likely dislodges a miniscule piece of frost from the inside of the bottle lid, providing that missing nucleation site. The ice crystals then grow on one another in a matter of seconds.

How solid the water freezes depends on the outside temperature, which isn't shared in the Reddit post. According to Physics Van, since water releases 80 calories per gram in heat when it freezes, typical kitchen freezer temperatures (which aren't more than a few degrees below freezing) usually aren't enough to freeze a bottle of water rock solid. The water that is freezing actually gives off enough heat to keep the rest of the water in a liquid state, resulting in slush. However, in cold-enough temperatures, a water bottle could indeed snap freeze quite solidly: Water can be supercooled down to as low as negative 43.6 degrees F (negative 42 degrees C), according to Physics Van.   

The liquid-to-solid transition isn't the only time that water behaves strangely. It can happen at water's boiling point, too. Water can become superheated when its temperature goes above 212 degrees F (100 degrees C) without transitioning to a gaseous state. Just as with supercooling, this happens in water pure enough and containers smooth enough not to provide nucleation sites. Just like ice crystals in supercooled water, the bubbles that start the boiling process have nothing to cling to. But superheated liquids can be quite a bit more dangerous than supercooled ones: Instead of turning to ice when moved, they can suddenly erupt into a gas, leading to an explosion of very hot liquid. Microwaving distilled water can cause this sort of eruption — a phenomenon you can watch on YouTube without any danger of second-degree burns.

Original article on Live Science. 

As a series of goopy platforms climb down a bolt in a mesmerizing GIF posted on Reddit, it almost looks as if Mario should hop from one to another.

But this isn't 1990's video-game graphics, it's real life. The GIF shows a demonstration of ferrofluid, a suspension of nanosize magnetic particles in oil. The magnetic particles are small and coated in a surfactant, which is a substance like soap that helps to keep the particles evenly distributed throughout the fluid, even when they're put next to a strong magnet, said Brandon Jackson, a doctoral candidate in mechanical engineering at Michigan Technological University, who has studied applications for ferrofluids.

In the GIF posted on Reddit, and in many other similar demonstrations seen on YouTube, Twitter and GIPHY, a magnet under the bolt provides the magnetic field. The liquid aligns itself with the invisible lines of the magnetic field, resulting in the spiky look. Surface tension holds the liquid together. Meanwhile, gravity pulls the liquid down the screw, resulting in the downward motion of the "platforms." [See More Fun Chemical Reaction GIFs]

The shape "is a minimum-energy solution between the gravitational energy, the surface tension energy and the magnetic energy," meaning the substance chooses the route that demands the least amount of energy, Jackson told Live Science.  

In some GIFs, the demonstrator holds additional magnets, which can alter the shape of the ferrofluid further, or move it around a surface like a snuffling hedgehog.

Actually, it's rocket science

NASA scientists invented ferrofluids in the 1960s when they were trying to figure out how to efficiently move rocket fuel from a tank into an engine in a zero-gravity environment. The researchers thought that by magnetizing the fuel with tiny iron oxide particles, they might be able to use a magnetic field to suck the fuel into the engine, leaving behind any pesky gas bubbles that could cause damage, according to NASA's Technology Transfer Program.

via GIPHY

Solid-rocket propellants obviated the need for the ferrofluids in space, but scientists quickly realized that ferrofluids could also be used to form seals to protect semiconductor chips during fabrication. Probably the most common industrial application today is in speakers, Jackson said. Ferrofluids are used to dampen vibrations in speaker components to prevent a distorted sound.

The future of ferrofluids

Ferrofluids make for fun science-fair demos, but they're also the subject of active research. One potential application, Jackson said, is in medicine. Some scientists have considered ferrofluids as an alternative to radiation treatment for cancer, he said. The fluids could be injected into tumors and then vibrated with a rapidly alternating magnetic field that would heat them, essentially "cooking" the tumor from the inside. Radiation similarly kills tumor cells with heat, but causes damage as it passes through tissue on the way to the cancer. Ferrofluids could similarly be used to target drugs to certain tissues, Jackson said.

Jackson and his colleagues have a different use for ferrofluids in mind. They're studying ferrofluids as self-assembling thrusters on tiny satellites. Traditional propulsion systems work well on large satellites, Jackson said, but an increasing number of the satellites shot into orbit are the size of cellphones or shoe boxes. Many use electrospray thrusters, which use tiny electrified needles to spray jets of fluid in order to propel the satellite. But the spiky hedgehog shape that ferrofluids form under the influence of a magnetic field can also shoot jets of ions — a form of propulsion that requires only a magnet, not a precision-fabricated needle.

Original article on Live Science. 

A single drop of solvent swirls like a tiny dancer atop a beaker of water, gradually jettisoning away little round bits of itself until nothing remains. Some who saw it thought it looked like a spinning galaxy, or the world's tiniest hurricane. All who saw it wondered what the heck was going on — and that includes the researchers who conducted the experiment in 2011.

The magical drop of solvent stars in a GIF called "A drop of dichloromethane on water spiraling out of existence as it evaporates," posted Thursday (Jan. 11) to Reddit's r/chemicalreactiongifs forum. Despite its newfound fame (more than 20 thousand upvotes within the first 24 hours), the GIF originated from a 2011 paper published in the German chemistry journal Angewandte Chemie.

The paper's thesis was simple: when you release a droplet of dichloromethane (DCM) solvent into a beaker of soapy water, it looks really, really cool. [Album: Prize-Winning Photos Taken Through a Microscope]

“This is a very easy experiment, and a very complicated phenomenon,” said Oliver Steinbock, a professor of chemistry at Florida State University, and the senior author on the study. "We were very surprised by it — and we still are.”

To set up the experiment, Steinbock and his fellow researchers filled several beakers with various concentrations of water and a common lab disinfectant called CTAB. Using a pipette, they added a single drop of DCM — a colorless liquid sometimes used as a degreaser — to each beaker, and filmed the results. Each trial took about 20-30 seconds total, and was visible with the naked eye.

So, what's going on here?

Each drop of DCM, which has a relatively low boiling point, began evaporating as soon as it left the pipette. But the surprises began when the droplets touched the soapy water solution.

"DCM has a higher density than water, so you'd expect it to sink right away," Steinbock told Live Science. "But instead, as soon as it touches the water, part of it spreads out and creates this sort of film that holds the droplet on the surface of the water… it’s like a boat that holds the droplet afloat.” (Though the DCM film is not visible in the viral GIF above, you can see it clearly in several other videos of the experiment that Steinbock posted to YouTube.)

Despite this boat-like film, a small part of the droplet does begin to sink. It's not visible from the top-down vantage point of this GIF; however, a tiny jet of falling bubbles forms underneath the droplet once it touches the water. The falling DCM jet slowly shrinks the droplet's volume, but also causes it to spin. "It’s a little bit like when you flush a toilet," Steinbock said. "The water has a tendency to start to rotate and twist. And that triggers the rotation of the drop that we begin to see."

Within a few seconds, the droplet is at once floating, rotating and evaporating. As a result of these combined forces, smaller droplets eventually start disengaging from the edge of the larger droplet. But instead of sinking themselves, they shoot out radially, moving straight ahead over the surface of the film until they, themselves, evaporate. 

"These droplets are self-propelled," Steinbock said. This is due to a phenomenon called the Marangoni effect, which states that a liquid with a high surface tension will pull more strongly than a liquid with a low surface tension. This difference in tension creates a force on the system that can lead to motion.

As the DCM in the experiment begins to evaporate, the droplet's surface tension lowers from the outside in. Smaller droplets begin to form at the large droplet's edge, until the relatively high surface tension of the surrounding water pulls the small droplets away in what Steinbock calls a "ballistic" trajectory. Each individual droplet moves straight ahead until its surface tension becomes equally unstable, leading to further fragmentation. Eventually, the droplets split so many times they can no longer be seen. (A 2017 paper in Physical Review Letters explains the phenomenon further.)

These and other forces continue spinning and shrinking the large DCM droplet until, suddenly, it loses its symmetry and sputters madly into total evaporation. Why the system suddenly goes from a state of apparent symmetry to total entropic chaos baffles even Steinbock and his fellow researchers. Across half a dozen experiments, they were unable to re-create the exact patterns seen in this GIF. "I was a little discouraged to understand how complicated it really is," Steinbock said.

However complicated, this little drop of solvent nevertheless spoke to something intrinsic and pure in many who saw it. As Reddit user MurderSlinky put it: "Never before have I related so much to a gif as I have to this tiny, insignificant dot of liquid spinning aimlessly in an endless, indifferent sea while slowly becoming nothing."

Originally published on Live Science.

Here is the science behind the coolest chemical reactions that you’ve seen on Reddit, YouTube or elsewhere on the web.

A glass tube filled with lightning sparked plenty of attention after it was posted on Reddit this month, but what's causing the dancing electrified bolts that seem to have no source? Turns out, the seemingly sci-fi contraption is a demonstration of the effect of electric fields and the same principles that make neon lights glow.

The video doesn't show it clearly, but the holder of the tube is standing near a Tesla coil, which is essentially two open circuits connected to a "spark gap," or a gap between two electrodes that creates a spark. The coil can generate an electric field that can have thousands of volts but very little current, which makes it safer than, for example, touching a wall socket, which has a lot of current but little voltage. That's because a large electric current — the amount of energy traveling in an electrical flow — can be dangerous even if the voltage (the pressure that pushes the electrons along a current) is low. But a small current is much less dangerous even under high voltage.

Carl Willis, a nuclear engineer in New Mexico who builds similar art projects, told Live Science in an email that one type of Tesla coil can produce an electric field around it that changes direction tens or hundreds of thousands of times per second. If you put a tube full of gas at low pressure near it, the electric field will excite the atoms of gas, stripping electrons from them. The result? The positively charged ions (those stripped atoms of gas) and electrons (which are negatively charged) move to opposite sides of the tube. [Nikola Tesla vs. Thomas Edison: Who Was the Better Inventor?]

Initially, this generates the sparks we see because the ionization process — as the gas gets stripped of its electrons — produces light. Ordinarily, the light wouldn't last long because the ions and electrons would all gather at either end of the tube, and that would be that. But when the electric field reverses direction, they don't have time to settle into their new positions. Instead, some of the electrons and ions recombine to form neutral atoms, which get ionized again as the now-reversed field drives the charged particles to change direction.

The gas inside the Reddit user's tube seems to be xenon, and the pressure inside the tube is on the order of a pound per square inch, Willis noted. (Xenon is characteristically bluish.)

How does this connect to neon lights?

 If, instead of putting the tube in an electric one connected  two electrodes to it — one at either end — the gas could be charged with an ordinary current source, like a wall socket, rather than a Tesla coil. This is exactly what happens in neon signs and fluorescent lights. (The latter appear white because the inside of the bulb is coated with a material that glows white in the presence of UV light, which is what gets generated inside the bulb.) Electrons from the current source smack into the electrons surrounding atoms of gas, exciting them and resulting in the release of a photon — light. Neon's characteristic wavelength is reddish.

 In typical fluorescent bulbs, mercury gas is mixed with argon, krypton or xenon, depending on the brand. (Xenon lights are popular in car headlights, for example.)  In that way, the current from the electrodes is serving the function that the electric field does for the xenon-filled tube in the video. In fact, one can get both fluorescent and neon lights to light up in the presence of an electric field – as demonstrated by Florian Dussopt Design Studio's "EM Table," which generates a small electromagnetic field around it. 

One could do this at home — Tesla coils are widely available, and a number of artists sell discharge tubes and electric-field setups here.

Originally published on Live Science.

That scratchy steel wool that cleans up your grimy pans is more than hardworking; it is absolutely magnificent when lit on fire, as Reddit user ChazDodge showed in a recent video that makes the wiry, burning puff look like the death of the planet Krypton.

Though it's not an explosion caused by a nuclear chain reaction — à la Krypton — the light show created by the burning steel wool results from high-speed oxidation.

Here's how it works: Anytime something burns, you're seeing oxidation. That means an atom, molecule or ion loses one or more electrons. Rust, for example, occurs when oxygen hits iron, and in the process the iron loses electrons and forms iron oxide. Rusting is a slow version of the reaction seen in the Reddit post of the burning (oxidizing) metal strips that make up steel wool. [Can Humans Spontaneously Combust?]

Yet, we use our stainless steel (which contains iron) cooking utensils without expecting them to burst into flame from an errant spark. What gives?

The reason a block of iron like a utensil doesn't catch fire is that the surface area is small, relative to the volume, Jason Benedict, an associate professor of chemistry at the University at Buffalo, told Live Science. Rusting iron actually generates some heat in the reaction, but it's a very small amount. In addition, a big block of iron can absorb and dissipate a lot of that heat energy before the block's temperature goes up. (You can see this effect in heating a metal spoon when stirring boiling pasta — a small one very quickly gets too hot to hold, while a bigger spoon takes longer.) 

Steel wool, on the other hand, is made of lots of thin strands, and so a lot more iron atoms are in contact with the oxygen in the air. When you add heat (as from a flame), you add energy to the iron, and that makes the iron more likely to react with other elements.

"When you're adding heat, you're overcoming an energy barrier to make the reaction happen faster," Benedict said. Once that reaction gets going, and because it generates heat itself, it heats neighboring atoms. In a block of iron, the heat gets dissipated to many other iron atoms. But in a thin fiber of iron, there's less solid material to absorb it (air absorbs heat, but much more efficiently than solids), so it keeps burning. The product of the burn is bits of rust, or iron oxide, just as the product of burning wood is black ash (or carbon).

Contact with oxygen is crucial to how fast and how hot the iron in steel wool burns — a pure-oxygen environment makes the flames a lot hotter, and the iron burns faster. (While steel wool is often covered in other chemicals — powdered soap, for example — only the iron is burning and mixing with oxygen.)   

Air is only 20 percent or so oxygen, so the burning happens at a kind of half-speed that looks like a cartoon dynamite fuse. That's what's happening in the video — there's enough oxygen to burn the iron, but not enough to get it to burst into flame all at once. Again, one can draw an analogy with wood: Blow on a small flame and the extra oxygen can make the wood burn faster, while if you close the vents on an old-fashioned wood stove, the fire dies down to glowing embers and burns more slowly.  

This is also why powdered metals burn easily and so are used in welding. Thermite is a good example — thermite is a mix of iron and aluminum powder that when heated enough will start reacting with oxygen and burning at a high temperature — enough to melt metal and weld. Thermite also shows up on the Fourth of July — it's an ingredient in the stuff that coats sparklers. 

Editor's Note: This article was updated to indicate that oxidation is the loss of electrons, not the gaining of oxygen as had been stated previously.

Originally published on Live Science.

Whoever said chemistry is boring hasn't seen YouTube user Amazing Timelapse's video showing a calculator melting into a surreal shape, reminiscent of a Salvador Dalí painting. Surprisingly, the calculator isn't melting at all, or even being heated.

It's literally dissolving.

In the video, Amazing Timelapse exposes the plastic calculator to acetone vapor, which can dissolve many types of plastics. Like other organic polymers, plastics are made of long-chain carbon-based molecules. That makes them incapable of dissolving in water, but vulnerable to organic solvents like acetone.

Water is a polar molecule, with a positive and negative end. And so when you plunk compounds that are also made of polar molecules — salt (sodium chloride), for instance — into water, the negatively charged parts of water (oxygen) tug on the positive molecules in the salt (sodium), while water's positively charged hydrogen molecules attract the salt's negative molecules (chloride). The result? The sodium and chloride get separated; they dissolve in the water. [More Chemistry Fun: How to Make Puffy Slime]

Plastics are different. The long carbon chains aren't polar — they don't have the same positive and negative sides. So water just bounces off the molecules and doesn't separate them from their fellows.

Acetone behaves differently with plastic, because acetone — which has the chemical formula (CH3)2CO — is not a polar molecule. That gives acetone a special power in relation to some polymers: It can knock them around and make the individual molecules in the polymer loosen their grip on their neighbors. According to the Open3DP website run by the University of Washington, plastics such as ABS (which legos are made of) or polyvinyl chloride (a component of piping) will start to lose their shape and cohesion. This is also the reason acetone is such a great superglue remover.

You could dissolve polymers in acetone like dropping sugar into coffee, or expose it to the vapor. Do that long enough, and the plastic on the outside of the calculator starts to absorb the acetone and effectively melt, creating a surreal and hypnotic sight. Only the metal and silicon components of the calculator remain intact.

Amazing Timelapse, who wouldn't reveal his or her identity (or identities), chose the vapor route. To get the weird effect, they suspended the calculator inside a container; the background was just a green screen, they said. They aren't professional photographers or chemists, they said, adding that the dissolution experiment was just aesthetically interesting.

Originally published on Live Science.

When the seasons change, they typically usher in new splashes of color — from the bright hues of blooming flowers to the rich tones of leafy trees. Now, a mesmerizing new video shows that these explosions of color are just as spectacular (and perhaps even more so) at teeny-tiny scales in the microscopic world.

In "Seasons," a video from Beauty of Science, an educational outreach organization, microscopic and macro-lens footage from previous projects were used to create a colorful display of the world's changing seasons. From crystals slowly growing to water droplets moving across the screen, the video clips are used to get viewers the feelings of the seasons as they change.

Yan Liang, a chemist and producer for Beauty of Science, said the organization's videos are "explorations of visual art under a microscope." [Small World: Gallery of Microscopic Beauty]

One previous Beauty of Science video series used in "Seasons" was Miniglobelet, which looked at crystallization, and color pigments interacting with various liquids (alcohol, water, ink remover, soft drinks, etc.) to create miniature dream-like sceneries.

Another video series included in "Seasons" was called Envisioning Chemistry, a project that captures chemical reactions using special photography techniques. Liang explained that they film reactions like the movement of metal, or the combustion of elements. Crystallization is a common reaction, with the videos capturing the creation of the crtystalline veins.

While crystallization, for example, has a clear connection to the winter season, the Beauty of Science team's use of colorants and other materials allows for the short film's seasonal transitions.

In other videos, Beauty of Science has tackled different experiments, including a recent film that shows how an M&M candy dissolves in water.

Capturing the microscopic world on video has long been an aspect of scientific research, but it has also been expanded to educational and entertainment purposes. The American Chemical Society, for example, has a video series on YouTube that explains different chemical reactions, from sparklers to fluorescence.

Original article on Live Science.

For the first time, scientists have accomplished a feat long thought impossible — they have measured the energy of incredibly short-lived arrangements of atoms that occur as chemical reactions are happening.

This finding could help shed light on the precise inner workings of chemical reactions too complex to understand by other methods, the researchers said.

The chemical reactions responsible for life, death and everything in between involve molecules transforming from one kind to another — essentially, from reactants to products. As chemical reactions occur, fleeting and unstable arrangements of atoms, known as transition states, exist as molecular bonds break and form between atoms. [Wacky Physics: The Coolest Little Particles in Nature]

"Reactants and products are stable valleys on either side of a mountain range, and the transition state is the pass," study lead author Joshua Baraban, a physical chemist at the University of Colorado Boulder, said in a statement. "Because it only exists as you go from one thing to another, it's never really been thought of as something that you can easily study directly."

Now, for the first time, scientists have measured the amount of energy required to reach a transition state.

"This is something that, if you asked people with a Ph.D. in chemistry, they'd say it was not possible to do," Baraban told Live Science. "There are textbooks that say this is not possible to do."

The researchers investigated a kind of chemical reaction known as an isomerization, in which a molecule undergoes a change of structure. They focused on a molecule known as acetylene, which consists of two carbon atoms and two hydrogen atoms.

When acetylene absorbs energy, there are two conformations it can adopt, which can be visualized by imagining the atoms as balls and the molecular bonds connecting the atoms as sticks. In acetylene, the carbon atoms are bound to each other and make up the middle of the molecule, and each carbon atom has one hydrogen atom attached to it.

One conformation has a zigzag shape, in which one hydrogen atom is positioned on one side of the carbon-carbon bond, while the other is on the other side of the carbon-carbon bond. The other conformation is shaped like a "U," with both hydrogen atoms on the same side of the carbon-carbon bond.

With a bit of energy, the zigzag version of acetylene can become the U-shaped kind, the researchers said. In between, a transitional state occurs where one of the hydrogen atoms is not positioned on either side of the carbon-carbon bond, but instead is almost in line with it.

The researchers used lasers to monitor changes in acetylene vibrations as the researchers gave more energy to the molecules. When specific levels of energy were reached, the patterns of vibrations changed in the kinds of ways expected near the transition state, the researchers said.

This kind of change in conformation is also an important part of how the eye works. "When light enters the eye, it causes this kind of change we see in acetylene, which starts a chain reaction that sends information that the eye has seen a photon to the brain," Baraban said.

The scientists also showed that they can use their technique to accurately predict the structure and energy of the transition state between hydrogen cyanide and hydrogen isocyanide. In hydrogen cyanide, a hydrogen atom is connected to a carbon atom, which, in turn, is bound to a nitrogen atom. In hydrogen isocyanide, a hydrogen atom is connected to a nitrogen atom, which, in turn, is bound to a carbon atom. The transition state between these molecules has one hydrogen atom, one carbon atom and one nitrogen atom bound to one another like a triangle.

Future research can analyze more complex reactions, such as ones where two molecules come together or one molecule breaks into two, the scientists said.

Baraban, along with study senior author Robert Field at MIT and colleagues, detailed their findings online today (Dec. 10) in the journal Science.

Follow Live Science @livescience, Facebook & Google+. Original article on Live Science.

This article was originally published at The Conversation. The publication contributed the article to Live Science's Expert Voices: Op-Ed & Insights.

If you stop and think about it for a moment, you will realise what an astonishing feat of precision engineering your colour printer is. It can take the primary colours – cyan, yellow, magenta and black – and mix them together carefully enough to achieve more than a million different hues and shades. Not only that but the drops of colour are mere nanolitres (billionths of a litre) in volume, each of which is then placed on the paper – assuming its not jammed in the feeder tray – with better than pinpoint accuracy.

Now a group of enterprising chemists from Tsinghua University are exploiting that precision engineering, which normally results in high-resolution colour prints, to screen millions of different chemical reactions. Their results have been published in the journal Chemical Communications.

Yifei Zhang and colleagues have been trying to understand reaction pathways in living things. Every chemical process that goes on in living organisms is controlled by a cascade of reactions. The steps in a cascade are mediated by protein molecules called enzymes. Each enzyme makes a small chemical alteration, like workers on a production line, to a molecule before passing its product onto the next enzyme. In this way, for example, plants build sugars from carbon dioxide and your food gets broken down and then reconstructed into other useful chemicals for your body.

The problem is that to understand these complicated processes by reconstructing them outside of a living cell is difficult. The concentrations of an enzyme relative to the next in the line is key. Get this wrong and bottle necks are formed in the production line, as one enzyme works faster than the next.

To figure out what are the right conditions to replicate a living cell’s workings, chemists must set up and monitor a vast number of reactions. Screening large numbers of reactions like this is often done using “96-well plates”, which are 96 tiny containers with a unique combination of chemicals in each. These reactions might be set up manually or, if the lab is well-funded, by an expensive robot. But even with the best robots available it can still be a slow process.

Colour printers are a lot cheaper than robots. And if the inks are replaced by solutions of enzymes then suddenly you have a device that has the potential to dispense more than a million different reaction mixtures.

That is just want Yifei and colleagues have done. Their printers were loaded with a series of enzymes that, when they work together in the correct ratios, produce coloured reaction products. These were printed directly onto paper where it was immediately obvious, from the intensity of a coloured dot, which reaction mixtures worked best.

In the test cases reactions were deliberately chosen that resulted in colour changes. This made for a nice quick visual indication of whether the system worked well. So for example one test started with glucose and a chemical called ABTS in the magenta cartridge, then the enzymes glucose oxidase (GOx) and horse-radish peroxidase (HRP) in the yellow and cyan cartridges. When they are mixed together the GOx removes a hydrogen from the glucose and adds it to oxygen, producing hydrogen peroxide. Next the HRP reacts this with the ABTS, which results in a green chemical.

The potential applications for these printer-based mixtures extend beyond curiosity-driven research on biological pathways. Yifei and colleagues have already shown that by loading the printer cartridges with the right enzymes they can use the set up to indicate the presence of glucose in a sample. Glucose in urine is a indication of diabetes, so their printer-based chemistry already has the potential to diagnose diabetes.

The result then could be a future where a trip to the doctors results in a printout of, quite literally, your urine and some enzymes alongside, after 30 seconds or so, a diagnosis and the prescription.

Mark Lorch does not work for, consult to, own shares in or receive funding from any company or organisation that would benefit from this article, and has no relevant affiliations.

This article was originally published on The Conversation. Read the original article. Follow all of the Expert Voices issues and debates — and become part of the discussion — on Facebook, Twitter and Google +. The views expressed are those of the author and do not necessarily reflect the views of the publisher. This version of the article was originally published on Live Science.

Celebrating in style

Scientists regularly watch spectacular microscopic bursts of activity happening inside us and other organisms. Using imaging techniques, including colorful dyes and graphic design programs, they generate pictures that could compete for the "oohs and ahhs" at any fireworks show. The grand finale: a better understanding of fundamental life processes that contribute to health and disease.

Here are just a few glimpses into cells captured by scientists in the course of their research funded by the National Institutes of Health. To see more fireworks on a microscopic scale, visit the online gallery for “Life: Magnified,” an exhibit of scientific images on display through November 2014 at Washington Dulles International Airport.

Liver Cell

Hepatocytes, like the one shown here, are the most abundant type of cell in the human liver. They play an important role in building proteins, producing bile (a liquid that aids in digesting fats) and chemically processing naturally occurring molecules like hormones as well as foreign substances like medicines and alcohol.

Nerve Cells in the Eye

Here we see the many layers of nerve cells in the retina of a ground squirrel. The top layer (green) is made up of cells called photoreceptors. These cells convert light into electrical signals that travel to the brain. The two best-known types of photoreceptors are rod and cone cells. Rods help us see under low-light conditions, and cones allow us to see vibrant colors in daylight.

Muscle Fibers

Of the three chicken muscle fibers shown here, the ones on the right and left are normal. The middle fiber is deficient in a protein called nebulin, which appears blue in the other fibers. Nebulin is critical to the structure and function of muscles, and its absence is associated with certain neuromuscular disorders.

Skin Cell

Skin Cell
This normal human skin cell has been treated with a natural chemical that triggers the formation of specialized protein structures that enable the cell to move. We depend on cell movement for such basic functions as wound healing and launching an immune response.

Fibroblasts

Fibroblasts
The cells shown here are fibroblasts, one of the most common cells in mammalian connective tissue. These particular cells came from a mouse. Scientists used them to test the power of a new microscopy technique that offers vivid views inside a cell. The mitochondria (green), cellular skeleton (red) and DNA within the nucleus (blue) are clearly visible.

Cells Lining Blood Vessel Walls 


This image captures the structure of a human endothelium, the thin layer of cells that lines our arteries and veins. The endothelium is a gatekeeper, controlling the movement of materials into and out of the bloodstream. Endothelial cells are held tightly together by specialized proteins that function like strong ropes (red) and others that act like cement (blue).

This Inside Life Science article was provided to LiveScience in cooperation with the National Institute of General Medical Sciences, part of the National Institutes of Health.

The Nobel Prize in chemistry has been awarded jointly to three scientists for laying the foundation for powerful computer models that are used to understand and predict complex chemical processes.

Martin Karplus, Michael Levitt and Arieh Warshel won the Nobel in chemistry "for the development of multiscale models for complex chemical systems," according to a statement by the Royal Swedish Academy of Sciences today (Oct. 9).

The trio's work was notable because they were able to apply both classical physics (that laid out by Isaac Newton) and quantum physics to their models of chemical processes and reactions. For instance, the quantum calculations were performed on the atoms and electrons within larger molecules, while classical calculations could simulate the reactions of these larger molecules.

Beginning in the 1970s, Warshel and Karplus worked to develop a computer model of retinal, a molecule in the eye's retina that has "free electrons," or those that can hop between atomic nuclei. While their model could handle both quantum and classical physics to simulate retinal, it could only simulate those molecules with mirror symmetry. That's where Levitt came in. Levitt and Warshel worked for several years, overcoming many obstacles, in their quest to develop a program that would allow quantum and classic theory to work side-by-side in a computer model of any type of molecule. [Photos: Stunning Peek Inside Molecules]

"Today the computer is just as important a tool for chemists as the test tube," according to a statement by the Academy. "Simulations are so realistic that they predict the outcome of traditional experiments."

Karplus, who was born in 1930 in Vienna, is a U.S. and Austrian citizen. He received his doctoral degree in 1953 from Cal Tech and is now at Université de Strasbourg, France, and Harvard University.

Levitt, who was born in 1947 in Pretoria, South Africa, is a U.S., British and Israeli citizen. He received a doctoral degree in 1971 from the University of Cambridge in the UK and is now at Stanford University School of Medicine.

Warshel, who was born in 1940 in Kibbutz Sde-Nahum, Israel, is a U.S. and Israeli citizen. He received his doctoral degree in 1969 from Weizmann Institute of Science in Israel and is now at University of Southern California, Los Angeles.

The three Laureates will share equally the award of $1.25 million (8 million Swedish krona).

Follow Jeanna Bryner on Twitter and Google+. Follow us @livescience, Facebook & Google+. Original article on LiveScience.