Periodic Table of Elements: Los Alamos National Laboratory 7, 6, 5, 4, 3, 0 History Plutonium is the second transuranium element of the actinide series. Element 93 was discovered in 1940/41 by Glenn T. Seaborg , Edwin M. McMillan , J. W. Kennedy , and A. C. Wahl by deuteron bombardment of uranium-238 in the 60-inch cyclotron at the University of California, Berkeley Lab. They first synthesized neptunium-238 (half-life 2.1 days) which subsequently beta-decayed to form a new heavier element with atomic number 94 and atomic weight 238 (half-life 87.7 years). It was fitting that element 94 be named after the next planetoid, Pluto following the precedence that uranium was named after the planet Uranus and neptunium after the planet Neptune.

Seaborg submitted a paper to the journal Physical Review in March 1941 documenting the discovery, but the paper was quickly withdrawn when it was found that an isotope of plutonium, Pu-239 could undergo nuclear fission making it useful in developing an atomic bomb. Pu-239 had a fission cross-section 50% greater than that of 235U, the best fissioning element known at that time. Seaborg was called away from Berkeley to lead the Plutonium Production Lab or "Met Lab" at the University of Chicago. The Met Lab was to produce useful quantities of plutonium as part of the secret Manhattan Project during World War II to develop an atomic bomb. On August 18, 1942, a trace quantity of plutonium was isolated and measured at the Met Lab for the first time. About 50 micrograms of Pu-239 combined with uranium and fission products was produced and only about 1 microgram was isolated. This was enough material for chemists to determine the new element's atomic weight. In November 1943 a few milligrams of PuF3 was reduced to create the first sample of plutonium metal.

Enough plutonium was produced to make it the first man-made element to be visible to the unaided eye. The nuclear properties of plutonium-239 were also being studied and researchers found that when hit with a neutron it fissions by releasing energy and more neutrons. These neutrons can hit neighboring atoms of Pu-239 and so on, in an exponentially fast chain-reaction, releasing a tremendous amount of energy. This energy could result in an explosion large enough to destroy a city or fuel a nuclear reactor. During WW II the three primary research and production sites of the Manhattan Project were the Plutonium Production Facility at what is now the Hanford Site , Washington, the Uranium Enrichment facilities at Oak Ridge, Tennessee , and the weapons research and design laboratory, now known as Los Alamos National Laboratory . In 1943, the first production reactor that made Pu-239 was the X-10 Graphite Reactor built at a facility in Oak Ridge, Tennessee that later became the Oak Ridge National Laboratory. The Manhattan Project produced the plutonium for the "Trinity Test" conducted in New Mexico by Los Alamos Laboratory Director Robert Oppenheimer and Army General Leslie Groves.

The world’s first atomic bomb ("The Gadget") was exploded near Socorro, New Mexico on July 16, 1945, resulting in an explosion with an energy equivalent of approximately 20,000 tons of TNT . The first atomic bomb used in war had a uranium core and was dropped on Hiroshima, Japan on August 6, 1945. The second atomic bomb used had a plutonium core and was nicknamed " Fat Man " because of its round shape. It was used to destroy Nagasaki, Japan in August 9, 1945, which put an end to WW II. Publication of the discovery and the naming of the new element plutonium was delayed until a year after the end of World War II . Seaborg originally considered the name "plutium", but later thought that it did not sound as good as "plutonium." Later, during the Cold-War era, large stockpiles of weapons-grade plutonium were built up by both the Soviet Union and the United States. Each year about 20 tons of plutonium is still produced as a by-product of the nuclear power industry.

As of 2007 it was estimated that the plutonium stockpile was about 500 tons, world-wide. Since the end of the Cold War these stockpiles have become a focus of nuclear proliferation concerns. In 2000, the United States and the Russian Federation mutually agreed to each dispose of 34 tons of weapon grade plutonium before the end of 2019 by converting it to a mixed uranium-plutonium oxide (MOX) fuel to be used in commercial nuclear power reactors. Today plutonium-239 remains an important component of nuclear weapons, and the United States maintains plutonium-related capabilities in support of national defense and global nuclear deterrence. Pu-239 for civilian nuclear power plants provides energy for many nations. Plutonium-238 continues to be vital to space exploration pushing the limits beyond which manned space exploration is possible and satisfying our quest for knowledge. Isotopes Twenty-three radioactive isotopes of plutonium have been characterized from mass numbers 228 to 247. Nine of these exhibit metastable states , though these all have half-lives less than one second.

The longest-lived isotopes are plutonium-244, with a half-life of 80.8 million years, plutonium-242, with a half-life of 373,300 years, and plutonium-239, with a half-life of 24,110 years. All of the remaining radioactive isotopes have half-lives less than 7,000 years. The primary decay modes of isotopes with mass numbers lower than plutonium-244, are spontaneous fission and α emission , mostly forming uranium and neptunium isotopes as decay products along with a variety of daughter fission products. The primary decay mode for isotopes with mass numbers higher than plutonium-244 is by β emission , mostly forming americium isotopes as daughter decay products. Plutonium-241 is the parent isotope of the neptunium decay series , decaying to americium-241 via β decay.

By far of greatest importance is the isotope 239Pu produced in extensive quantities in nuclear reactors from natural uranium: 238U (n, gamma) → 239U (beta) → 239Np (beta) → 239Pu Plutonium-238 with a half-life of 87.7 years is another important isotope. Both Pu-239 and Pu-238 have many practical applications as discussed below. Properties Plutonium is unique among the elements in its physicochemical complexities by virtue of its position at a transitional location in the periodic table where the 5f electrons are at the border between delocalized (not associated with a single atom) and localized (associated with a single atom) behavior and it is considered one of the most complex of the elements. Plutonium also sits near the juncture where the actinide series transitions from main d-block element chemistry to rare earth like behavior as a result of the actinide contraction. Because of its defense and commercial importance, plutonium is one of the most intensely investigated of elements.

Plutonium metal has a bright silvery appearance at first and takes on a dull gray, yellow or olive green tarnish when oxidized in air. A relatively large piece of plutonium is warm to the touch because of the energy given off by alpha decay. Larger pieces will produce enough heat to boil water. The metal readily dissolves in concentrated mineral acids. Plutonium metal normally has six allotropes or crystal structures; alpha (α), beta (β), gamma (γ), delta (δ), delta prime (δ') and epsilon (ε). It forms a seventh phase (zeta, ζ) under high temperature and a limited pressure range. These allotropes have very similar energy levels but significantly varying densities (from 16.00 to 19.86 grams/cm3) and crystal structures . This makes plutonium very sensitive to changes in temperature, pressure, or chemistry, and allows for dramatic volume changes following phase transitions . At room temperature plutonium is in its alpha (α) form , the most common structural form of the element. It is as hard and brittle as cast iron unless alloyed with other metals to form the room-temperature stabilized delta (δ) phase which makes it soft and ductile.

Unlike most metals, it is not a good conductor of heat or electricity . It has a low melting point (640 °C) and an unusually high boiling point (3,228 °C). A ring of 99.96% pure electro-refined plutonium (Los Alamos National Laboratory) Plutonium can form alloys and intermediate compounds with most other metals. Gallium, aluminum, americium, scandium and cerium can stabilize the δ phase of plutonium metal. Nuclear fuel pellets can be formed by alloying plutonium with various metals such as: aluminum; zirconium; cerium; cerium-cobalt; uranium-titanium, uranium-zirconium and uranium-molybdenum. Thorium-plutonium-uranium alloys were investigated as a nuclear fuel for fast breeder reactors. A plutonium-gallium-cobalt alloy (PuCoGa5) was found to be an unconventional superconductor, showing superconductivity below 18.5 Kelvin, an order of magnitude higher than the highest between heavy fermion systems known. Plutonium forms compounds with a variety of other elements. Plutonium reacts with pure hydrogen, forming plutonium hydride s.

It also reacts readily with oxygen, forming PuO and PuO2 as well as intermediate and sub-stoichiometric oxides. The metal reacts with the halogens , giving rise to trivalent Pu compounds with the general formula PuX3 where X can be F , Cl , Br or I and tetravalent plutonium compounds such as PuF4 . The following oxyhalides are observed: PuOCl, PuOBr and PuOI. Plutonium reacts with carbon to form PuC, nitrogen to form PuN and silicon to form PuSi2. Pu3+ and Pu4+ oxalates are important intermediates that are calcined to form oxides as a step in plutonium processing. Other important compounds in reprocessing are fluoride, peroxide, acetylacetone, carbonate and hydroxide. The color displayed by plutonium solutions depends on both the oxidation state and the extent of complexation by various ligands.

In aqueous solution plutonium exhibits five ionic valence states: Pu+3 (blue lavender), Pu+4 (salmon-colored, when uncomplexed), PuO+ (lavender), PuO+2 (orange-brown) and PuOxOHy (dark green in basic solution). The pentavalent ion, PuO+ is unstable in aqueous solutions and it disproportionates into Pu+4 and PuO+2. However, PuO2+ can be stabilized in aqueous solution in a narrow pH range around 4.5. By virtue of the close proximity of the electrode potentials of the various plutonium redox couples (~ 1 Volt/NHE), four oxidation states can co-exist in solution simultaneously: Pu3+, Pu4+, PuO2+ and PuO22+. Various oxidation states of plutonium in aqueous solution; all in non-complexing perchloric acid, except Pu(VII) in strong base. (D. E. Hobart and P.D. Palmer, Los Alamos National Lab.) Pu4+ is a "hard" (ionic) cation with the largest electronic charge of plutonium ions and it forms complexes with a variety of inorganic and organic ligands.

In dilute perchloric acid, Pu4+ is un-complexed and is salmon-colored. However in concentrated acids, Pu4+ forms anionic complexes such as: Pu(NO3)62- (dark green) and Pu(Cl)62- (brick red). Pu4+, having a high ionic charge readily hydrolyzes (combines with hydroxide ion) at near-neutral pH values forming a green colloidal suspension that behaves like a solution but is actually a solid precipitate that can be separated by ultra-centrifugation. Plutonium-organic complexes are very important for separation, reprocessing, and purification and include: Tributyl phosphate (TBP); Di-(2-ethylhexyl)phosphoric acid (DEHPA or HDEHP); octyl(phenyl)-N,N-diisobutyl-carbamoylmethylphosphine oxide (CMPO); crown-ethers; and many others. Production Plutonium is the heaviest primordial element in existence by virtue of its most stable isotope, plutonium-244, whose half-life of about 80 million years is just long enough for the element to be found in trace quantities in nature.

In 1971 mass spectrometric measurements of plutonium isolated from Precambrian bastnasite (a lanthanide chloride-fluoride mineral) by Hoffman and coworkers confirmed the presence of Pu-244 in nature. Although the existence of Pu-244 as an extinct radioactive element had been postulated to explain the xenon isotope ratios observed in meteorites, this is the first indication of its present existence in nature. Plutonium also exists in trace quantities in naturally occurring uranium ores. It is formed in much the same manner as neptunium: by irradiation of natural uranium with the neutrons followed by beta decay. Exceedingly small amounts of Pu-238, attributed to the extremely rare double-beta decay of U-238, have been found in natural uranium samples. Plutonium was most likely formed by neutron activation of natural U-238 at the Oklo natural reactor but if formed it has long since decayed away. Plutonium is for the most part a by-product of nuclear reactions in reactors where some of the neutrons released by the fission process convert U-238 nuclei into plutonium.

Plutonium-238 and -239 are the most widely synthesized isotopes. Plutonium-239 is synthesized via the following reaction using uranium (U) and neutrons (n) via beta decay (β−) with neptunium as an intermediate: Plutonium-238 is synthesized by bombarding uranium-238 with deuterons in the following first discovery reaction: Uses Plutonium has assumed the position of dominant importance among the transuranium elements because of its use as an explosive ingredient in nuclear weapons and the place which it holds as a key material in the development of industrial use of nuclear power. During fission, a fraction of the binding energy, which holds a nucleus together, is released as a large amount of electromagnetic and kinetic energy which is quickly converted to thermal energy. Fission of a kilogram of plutonium-239 can produce an explosion equivalent to 21,000 tons of TNT which is equivalent to about 22 million kilowatt hours of heat energy. In 1982 it was estimated that about 300,000 kg had accumulated.

The most common chemical process, PUREX (Plutonium–URanium EXtraction) reprocesses spent nuclear fuel to extract plutonium and uranium which can be used to form a mixed U/Pu oxide or "MOX" fuel for reuse in nuclear power reactors. MOX fuel production is also a good mechanism to reduce excessive defense plutonium stockpiles for peaceful purposes, which in effect is forging "swords into plowshares." Plutonium isotopes undergo radioactive decay, which produces decay heat . Different isotopes produce different amounts of heat per mass. Pu-238 with a half-life of 88 years has a relatively high heat production rate which makes it useful as a power source with a long service life. The decay heat is usually listed as watt/kilogram, or milliwatt/gram. Pu-238 is a heat source in radioisotope thermoelectric generators, which are used to power spacecraft and extra-terrestrial rovers. As a power and heat source, Pu-238 has also been used to power instruments left on the Moon by Apollo astronauts, weather satellites and interplanetary probes and powers the Cassini Saturn mission and the Mars rovers.

Plutonium-238 was at one time used successfully to power artificial heart pacemakers but has been replaced by lithium-based primary cells. Plutonium-238 was studied as a way to provide supplemental heat to scuba divers. Pu-238 mixed with beryllium is a convenient method to generate neutrons. Pu-238 powers the Cassini Mission now orbiting Saturn. (NASA) Hazards Plutonium, along with all of the other transuranium elements, is a radiological hazard and must be handled with specialized equipment and precautions. Animal studies have found that a few milligrams of plutonium per kilogram of tissue are lethal. Plutonium is more dangerous when inhaled than when ingested. When inhaled, plutonium can pass into the bloodstream moving throughout the body and into the bones, liver, or other body organs. Plutonium that reaches body organs generally stays in the body for decades and continues to expose the surrounding tissue to radiation and thus may cause cancer. Precautions must also be taken to prevent the unintentional formulation of a critical mass.

Pure Pu-239 metal present in sufficient quantity with an appropriate geometry can form a critical mass and pose a serious and often lethal criticality event. Low levels of plutonium in the environment is a result of world-wide fallout from numerous above-ground nuclear tests during the Cold War (now banned) pose little health hazard. Disposal of plutonium waste from nuclear power plants and dismantled nuclear weapons built during the Cold War is a nuclear-proliferation and environmental concern. Further Reading D. L. Clark, S. S. Hecker, G. D. Jarvinen, and M. P. Neu (2006) "Plutonium", Chapter 7 in The Chemistry of the Actinide and Transactinide Elements, Third Edition, L. R. Morss, J. Fuger, and N. M. Edelstein, Eds, Springer Publishers. Los Alamos Science (2000) Number 26, Volume I and II, Los Alamos National Laboratory report LA-UR-00-4100. Updated by Dr. David Hobart, August 2013.

LA-UR-13-26653 Facts About Plutonium Facts About Plutonium By Alina Bradford, Live Science Contributor | December 7, 2016 09:30pm ET MORE These buttons of refined plutonium metal were used in the core of the Fat Man bomb dropped on Nagasaki. Credit: U.S. Department of Energy Plutonium is a radioactive, silver metal that can be used to create or destroy. While it was used for destruction soon after it was made, today the element is used mostly for creating energy throughout the world. Plutonium was first produced and isolated in 1940 and was used to make the "Fat Man" atomic bomb that was dropped on Nagasaki at the end of World War II, just five years after it was first produced, said Amanda Simson, an assistant professor of chemical engineering at the University of New Haven.

Just the facts Here are the properties of plutonium, according to the Los Alamos National Laboratory : Atomic number: 94 Melting point: 1,184 F (640 C) Boiling point: 5,842 F (3,228 C) Discovery & history Plutonium was discovered in 1941 by scientists Joseph W. Kennedy, Glenn T. Seaborg, Edward M. McMillan and Arthur C. Wohl at the University of California, Berkley. The discovery occurred when the team bombarded uranium-238 with deuterons that had been accelerated in a cyclotron device, which created neptunium-238 and two free neutrons. The neptunium-238 then decayed into plutonium-238 through beta decay. This experiment wasn't shared with the rest of the scientific community until 1946, after World War II. Seaborg submitted a paper on their discovery to the journal Physical Review in March 1941, but the paper was removed when it was discovered that an isotope of plutonium, Pu-239, may be used to create an atomic bomb.

Soon, Seaborg was sent to lead the Plutonium Production Lab, also known as the Met Lab, at the University of Chicago, according to the Los Alamos National Laboratory. The purpose of the lab was to create plutonium as part of the Manhattan Project. The Manhattan Project was a secret venture during World War II that worked exclusively to develop an atomic bomb. On August 18, 1942, they had their first big success. They were able to create a trace quantity of plutonium that was visible to the eye. It equaled only around 1 microgram. From the tiny sample, scientist determined plutonium's atomic weight. The Manhattan Project eventually produced enough plutonium for the "Trinity Test." During the test, the world's first atomic bomb, or the "The Gadget," was exploded near Socorro, New Mexico, on July 16, 1945, by Los Alamos Laboratory director Robert Oppenheimer and Army General Leslie Groves. Of the test, Oppenheimer said, "We knew the world would not be the same. A few people laughed, a few people cried. Most people were silent.

I remembered the line from the Hindu scripture, the Bhagavad-Gita. Vishnu is trying to persuade the Prince that he should do his duty and to impress him takes on his multi-armed form and says, 'Now I am become Death, the destroyer of worlds.' I suppose we all thought that, one way or another," according to the Royal Society of Chemistry . The explosion had the energy equivalent of approximately 20,000 tons of TNT. The first war-use atomic bomb dropped on Hiroshima, Japan, on August 6, 1945. That atomic bomb, dubbed "Little Boy," had a uranium core, though. The second bomb, dropped on Nagasaki, Japan, in August 9, 1945, had a plutonium core. The "Fat Man," as it was called, hastened the end of World War II. Properties of plutonium Freshly prepared plutonium metal has a silvery bright color but takes on a dull gray, yellow, or olive green tarnish when oxidized in air. The metal quickly dissolves in concentrated mineral acids.

A large piece of plutonium feels warm to the touch because of the energy given off by alpha decay; larger pieces can produce enough heat to boil water. At room temperature alpha-form plutonium (the most common form) is as hard and brittle as cast iron. It can be alloyed with other metals to form the room-temperature stabilized delta form, which is soft and ductile. Unlike most metals, plutonium is not a good conductor of heat or electricity. It has a low melting point and an unusually high boiling point. Plutonium can form alloys and intermediate compounds with most other metals, and compounds with a variety of other elements. Some alloys have superconductive abilities and others are used to make nuclear fuel pellets. Its compounds come in a variety of colors, depending on the oxidation state and how complex various ligands are. In aqueous solution there are five valance ionic states. Plutonium, along with all of the other transuranium elements, is a radiological hazard and must be handled with specialized equipment and precautions. Animal studies have found that a few milligrams of plutonium per kilogram of tissue are lethal. Sources Plutonium generally isn't found in nature. Trace elements of plutonium are found in naturally occurring uranium ores.

Here, it is formed in a way similar to neptunium: by irradiation of natural uranium with neutrons followed by beta decay. Primarily, however, plutonium is a byproduct of the nuclear power industry. Each year, around 20 tons of plutonium is produced, according to the Los Alamos National Laboratory. Spent nuclear fuel can also be reprocessed to separate usable plutonium from other elements in the fuel. Atmospheric weapons testing in the 1950s and 1960s left tons of plutonium in the Earth's atmosphere that is still there today, according to the World Nuclear Association . Uses For the most part, plutonium isn't used for much. In fact, of the five common isotopes, only two of plutonium's isotopes, plutonium-238 and plutonium-239, are used for anything at all. Plutonium-238 is used to make electricity for space probes using radioisotope thermoelectric generators. These generators are switched on when the probes can't get enough solar power because they have traveled too far away from the sun. Some probes that use plutonium-238 are Cassini and Galileo.

When concentrated enough, plutonium-239 undergoes a fission chain reaction. Because of this, it is used in nuclear weapons and some nuclear reactors. In fact, one of the biggest uses for plutonium is energy. According to the World Nuclear Association, over one-third of the energy produced in most nuclear power plants comes from plutonium. Plutonium is the main fuel in fast neutron reactors. Who knew? For decades, scientists wondered why plutonium didn't act like other metals in its group. For instance, plutonium is a poor conductor of electricity and it doesn't stick to magnets. Now researchers have figured out where its "missing magnetism" has been hiding out and it has to do with the wacky behavior of the electrons in the element's outer shell. Unlike other metals, which have a set number of electrons in their outer shells, when in a ground state, plutonium can have four, five or six electrons there. This fluctuating number of outer-shell electrons explains why plutonium isn't magnetic: In order for an atom to interact with magnets the unpaired electrons in its outer shell must line up in a magnetic field.

[ Read more about plutonium's missing magnetism ] Plutonium's most stable isotope, plutonium-244, can last a long time. It has a half-life of about 82 million years and decays into uranium-240 through alpha decay, according to the Jefferson Lab . Plutonium was named after the planet, Pluto. This is because it came after Uranium, which was named after the planet Uranus, and neptunium, which was named after the planet Neptune. Plutonium emits neutrons, beta particles and gamma rays. Additional resources Uranium - Element information, properties and uses | Periodic Table Chemistry in its element: uranium (Promo) You're listening to Chemistry in its element brought to you by Chemistry World, the magazine of the Royal Society of Chemistry. (End promo) Chris Smith For Chemistry in its element this week, can you guess what connects boat keels, armour piercing weaponry, beautiful coloured glass that you can track down with a geiger counter and more oxidation states than a chemist can shake a glass rod at. If not, here's Polly Arnold with the answer.

Polly Arnold Uranium is certainly one of the most famous, or perhaps I should say infamous, elements. It is the heaviest naturally occurring element. It is actually more abundant in the earth's crust than silver. It is one of eight elements named in honour of celestial objects, but you might not think that uranium deserves to be named after the planet Uranus. The lustrous black powder that the chemist Klaproth isolated from the mineral pitchblende in 1789 - just eight years after Uranus was discovered - was in fact an oxide of uranium. Not until fifty two years later did Eugène Melchior Peligot reduced uranium tetrachloride with potassium, and from these harsher conditions obtained the pure silvery white metal at last. Samples of the metal tarnish rapidly in air, but if the metal is finely divided, it will burst into flames. Uranium sits amongst the actinides, the second shell of metals to fill their f-orbitals with valence electrons, making them large and weighty. Chemically, uranium is fascinating. Its nucleus is so full of protons and neutrons that it draws its core electron shells in close.

This means relativistic effects come into play that affect the electron orbital energies. The inner core s electrons move faster, and are drawn in to the heavy nucleus, shielding it better. So the outer valence orbitals are more shielded and expanded, and can form hybrid molecular orbitals that generated arguments over the precise ordering of bonding energies in the uranyl ion until as recently as this century. This means that a variety of orbitals can now be combined to make bonds, and from this, some very interesting compounds. In the absence of air, uranium can display a wide range of oxidation states, unlike the lanthanides just above it, and it forms many deeply coloured complexes in its lower oxidation states. The uranium tetrachloride that Peligot reduced is a beautiful grass-green colour, while the triiodide is midnight-blue. Because of this, some regard it as a 'big transition metal'. Most of these compounds are hard to make and characterise as they react so quickly with air and water, but there is still scope for big breakthroughs in this area of chemistry.

The ramifications of relativistic effects on the energies of the bonding electrons has generated much excitement for us synthetic chemists, but unfortunately many headaches for experimental and computational chemists who are trying to understand how better to deal with our nuclear waste legacy. In the environment, uranium invariably exists as a dioxide salt called the uranyl ion, in which it is tightly sandwiched between two oxygen atoms, in its highest oxidation state. Uranyl salts are notoriously unreactive at the oxygen atoms, and about half of all known uranium compounds contain this dioxo motif. One of the most interesting facets of this area of uranium chemistry has emerged in the last couple of years: A few research groups have found ways to stabilise the singly reduced uranyl ion, a fragment which was traditionally regarded as too unstable to isolate. This ion is now beginning to show reactivity at its oxygen atoms, and may be able to teach us much about uranium's more radioactive and more reactive man-made sisters, neptunium and plutonium - these are also present in nuclear waste, but difficult to work with in greater than milligram quantities. Outside the chemistry lab, uranium is best known for its role as a nuclear fuel.

It has been at the forefront of many chemists' consciousness over recent months due to the international debate on the role that nuclear power can play in a future as a low-carbon energy source, and whether our new generations of safer and efficient power stations are human-proof. To make the fuel that is used to power reactors to generate electricity, naturally occurring uranium, which is almost all U-238, is enriched with the isotope U-235 which is normally only present in about 0.7 %. The leftovers, called depleted uranium, or DU, have a much-reduced U-235 content of only about 0.2 %. This is 40 % less radioactive than natural uranium, and the material that we use to make compounds from in the lab. Because it is so dense, DU is also used in shielding, in the keels of boats and more controversially, in the noses of armour-piercing weapons. The metal has the desirable ability to self-sharpen as it pierces a target, rather than mushrooming upon impact the way conventional tungsten carbide tipped weapons do. Critics of DU weaponry claim it can accumulate around battlefields.

Because uranium is primarily an alpha-emitter, its radioactivity only really becomes a problem if it gets inside the body, where it can accumulate in the kidneys, causing damage. However, uranium is also a heavy metal, and its chemical toxicity is of greater importance - it is approximately as toxic as lead or mercury. But uranium doesn't deserve it's image as one of the periodic table's nasties. Much of the internal heat of the earth is considered to be due to the decay of natural uranium and thorium deposits. Perhaps those looking to improve the public image of nuclear power should demand the relabelling of geothermal ground-source heat pumps as nuclear? The reputation of this element would also be significantly better if only uranium glass was the element's most publicly known face. In the same way that lead salts are added to glass to make sparkling crystal glassware, uranyl salts give a very beautiful and translucent yellow-green colour to glass, although glassmakers have experimented to produce a wide range of gem-like colours. An archaeological dig near Naples in 1912 unearthed a small green mosaic tile dated back to 79 AD, which was reported to contain uranium, but these claims have not been verified.

However in the early-19th and early 20th century it was used widely in containers and wine-glasses. If you think that you own a piece, you can check with a Geiger counter, or by looking for the characteristic green fluorescence of the uranium when held under a UV-lamp. Pieces are generally regarded as safe to drink from, but you are advised not to drill holes in them, or wear them. Fair enough. Chris Smith Or inadvertently eating it too, presumably. That was Edinburgh University chemist Polly Arnold explaining the softer side of the armour piercing element Uranium. Next week Andrea Sella will be introducing us to some crystals with intriguing properties. Andrea Sella "It's amazing stuff. You HAVE to see this." He pulled out of his pocket a sample vial containing some stunning pink crystals that glinted alluringly. "Wow!" I said - you can always impress a chemist with nice crystalline products. "It gets better." he said mysteriously. He beckoned me into a hallway. "Look" he said. As the crystals caught the light from the new fluorescent lights hanging from the ceiling, the pink colour seemed to deepen and brighten up. "Wow!" I said again.

We moved the crystals back into the sunlight and the colour faded again, and moving the crystals back and forth they glowed and dimmed in magical fashion. Chris Smith But what did they contain? Well, the answer's Erbium and you can hear all about it in next week's Chemistry in its element. I'm Chris Smith, thank you for listening and goodbye. (Promo) Chemistry in its element is brought to you by the Royal Society of Chemistry and produced by thenakedscientists.com . There's more information and other episodes of Chemistry in its element on our website at chemistryworld.org/elements . (End promo) Periodic Table of Elements: Los Alamos National Laboratory 6, 5, 4, 3, 0 History The use of uranium in its natural oxide form dates back to 79 A.D. when it was used as a yellow coloring agent in ceramic glazes. Yellow glass with 1% uranium oxide was found in an ancient Roman villa near Naples, Italy. In the late Middle Ages, pitchblende was extracted from the silver mines and was used as a coloring agent in the glassmaking industry. The identification of uranium as an element is generally credited to Martin H. Klaproth .

While experimenting with pitchblende in 1789, he concluded that it contained a new element, which he named after the newly discovered planet Uranus (named after the Greek god of the sky or heaven). What Klaproth actually identified was not the pure element but uranium oxide. The pure metal was first isolated in 1841 by Eugène-Melchior Péligot, who reduced anhydrous uranium tetrachloride with potassium metal. In 1896 Antoine H. Becquerel discovered that uranium exhibited invisible light or rays; it was radioactivity. In 1934 research by Enrico Fermi and others eventually led to the use of uranium fission in the first nuclear weapon used in war and later in the peaceful use of uranium as fuel in nuclear power production. An ensuing arms race during the Cold War between the United States and the Soviet Union produced tens of thousands of nuclear weapons that used uranium metal and uranium-derived plutonium-239. The security of those weapons and their fissile material following the breakup of the Soviet Union in 1991 is an ongoing concern.

In 1972 French physicist Francis Perrin discovered ancient and no longer active prehistoric natural nuclear fission reactors in uranium ore deposits at the Oklo mine in Gabon, West Africa, collectively known as the Oklo Fossil Reactors. The ore deposit is 1.7 billion years old; at that time, uranium-235 constituted about 3% of the total uranium on Earth (0.72% today). This is high enough to permit a sustained nuclear fission chain reaction to occur, provided other supporting geologic conditions exist. Isotopes Uranium is weakly radioactive because all its isotopes are unstable. There are 27 known isotopes of uranium ranging in atomic weights 217–219, 222–240 and 242, with half-lives of from billions of years to a few nanoseconds. Naturally occurring uranium consists of three major isotopes: 238U (99.28% abundance), 235U (0.71%), and 234U (0.0054%).

(The US DOE has adopted the value of 0.711 as being their official percentage of 235U in natural uranium.) All three isotopes are radioactive, with small probabilities of undergoing spontaneous fission but preferentially decaying by alpha emission. The half-life of uranium-238 is about 4.47 billion years and that of uranium-235 is 704 million years, making them useful in dating the age of the Earth. It also suggests that half of the uranium that existed from the formation of the Earth has decayed to other radioactive elements and eventually to stable elements. Much of the internal heat of the earth is thought to be attributable to the decay of uranium and thorium radio-isotopes. Uranium-238 is an α-particle emitter (occasionally, it undergoes spontaneous fission), decaying through the "Uranium Series" of nuclear decay , which has 18 members, all of which eventually decay into lead-206 , by a variety of different decay paths. The decay series of 235U , which is called the actinium series has 15 members, all of which eventually decay into lead-207.

The constant rates of decay in these decay series make the comparison of the ratios of parent to daughter elements useful in radiometric dating . Uranium-234 is a member of the "Uranium Series" , and it decays to lead-206 through a series of relatively short-lived isotopes. Uranium-233 is made from thorium-232 by neutron bombardment, usually in a nuclear reactor, and 233U is also fissile. Its decay series ends with thallium-205. Sources Uranium is the heaviest naturally-occurring element available in large quantities. The heavier “transuranic” elements are either man-made or they exist only as trace quantities in uranium ore deposits as activation products. Uranium occurs naturally in low concentrations of a few parts per million in soil, rock and water, and is commercially extracted from uranium-bearing minerals. Uranium, not as rare as once thought, is now considered to be more plentiful than mercury , antimony , silver , or cadmium , and is about as abundant as molybdenum or arsenic .

It occurs in numerous natural minerals such as pitchblende, uraninite, carnotite, autunite, uranophane, and tobernite. It is also found in phosphate rocks, lignite, monazite sands, and is recovered commercially from these sources. The United States Department of Energy purchases uranium in the form of acceptable U3O8 concentrates. This incentive program has greatly increased the known uranium reserves. Properties Pure uranium is a silvery white, weakly radioactive metal, which is harder than most elements. It is malleable, ductile, slightly paramagnetic, strongly electropositive and is a poor electrical conductor. Uranium metal has very high density, being approximately 70% denser than lead, but slightly less dense than gold. Uranium metal exhibits in three crystallographic modifications: alpha --> (688°C) --> beta --> (776°C) --> gamma. Uranium is pyrophoric when finely divided. It is a little softer than steel and is attacked by cold water in a finely divided state.In air, uranium metal becomes coated with a layer of oxide.

Acids dissolve the metal, forming the +3 oxidation state which oxidizes rapidly by water and air to form higher oxidation states. Uranium metal is unaffected by alkalis. Uranium metal can be prepared by reducing uranium halides with alkali or alkaline earth metals or by reducing uranium oxides by calcium, aluminum, or carbon at high temperatures. The metal can also be produced by electrolysis of KUF5 or UF4, dissolved in a molten salt mixture of CaCl2 and NaCl. High-purity uranium can be prepared by the thermal decomposition of uranium halides on a hot filament. Uranium metal reacts with almost all nonmetallic elements and their compounds , with reactivity increasing with temperature. Hydrochloric and nitric acids dissolve uranium, but non-oxidizing acids other than hydrochloric acid attack the element very slowly. When finely divided, it can react with cold water. In air, uranium metal oxidizes and becomes coated with a dark layer of uranium oxide .

Uranium forms a variety of alloys and compounds with the most important oxidation states being uranium(IV) and uranium(VI), and their two corresponding oxides are, respectively, uranium dioxide , UO2 and uranium trioxide , UO3. Besides the oxides, other Important uranium compounds include fluorides, chlorides, bromides, iodides, carbonates, hydrides, carbides, nitrides, phosphates, etc. At room temperatures, uranium hexafluoride, UF6, has a high vapor pressure , making it useful in the gaseous diffusion process used to separate the rare U-235 from the common U-238 isotope. Uranium hydrides, nitrides and carbides are relatively inert semimetallic compounds that are minimally soluble in acids and have been used as stable fuel pellets in nuclear power reactor technology. Uranium exists in aqueous solutions in the +3, +4, +5, and +6 oxidation states. Oxidation state +6 as the UO22+ ion (yellow in color) is the most stable state in solution.

Uranium in the +5 state as the UO2+ ion is colorless, quite unstable and disproportionates (reacts with itself) to form the +6 and +4 states. The +4 state (green) is reasonably stable in solution, but the +3 state (dark green or dark red depending on the illumination source - daylight vs fluorescent light) is unstable and easily oxidizes to +4. The +4 state in near-neutral pH solutions readily hydrolyzes to form black oxy-hydroxide precipitates. Uses Uranium was used in as coloring agents in ceramic glazes and glass in ancient Rome and in the Middle Ages producing orange-red to lemon yellow hues. More recently it was used as an orange glaze in contemporary Fiestaware© dishware but was later discontinued for health reasons. Many contemporary uses of uranium exploit its unique nuclear properties. Uranium-235 has the distinction of being the only naturally occurring fissile isotope . This means it can be split into two or three fragments ( fission products ) by thermal neutrons.

Uranium-238 is fissionable by fast neutrons, and is fertile , meaning it can be transmuted to fissile plutonium-239 in a nuclear reactor . Another fissile isotope, uranium-233 , can be produced from natural thorium and is also important in nuclear technology. While uranium-238 has a small probability for spontaneous fission or even induced fission with fast neutrons, uranium-235 and to a lesser degree uranium-233 have a much higher fission cross-section for slow neutrons. In sufficient concentration, these isotopes maintain a sustained nuclear chain reaction . This generates the heat in nuclear power reactors, and produces the fissile material for nuclear weapons . This nuclear conversion can be brought about in breeder reactors where it is possible to produce more new fissionable material than the fissionable material used in maintaining the chain reaction. Depleted uranium (238U) (depleted of uranium-235) is used in balistic armor penetration and as armor plating .

Uranium-238 is not fissile, but is a fertile isotope, because after neutron activation it can produce plutonium-239 , another fissile isotope. Indeed, the238U nucleus can absorb one neutron to produce the radioactive isotope uranium-239 . 239U decays by beta emission to neptunium-239, also a beta-emitter, that decays in its turn, within a few days into plutonium-239. 239Pu was used as fissile material in the first atomic bomb detonated in the " Trinity test " on 15 July 1945 in New Mexico . Uranium-235 is of even greater importance because it is the key to utilizing uranium. 235U, while occurring in natural uranium to the extent of only 0.71%, is so fissionable with slow neutrons that a self-sustaining fission chain reaction can be made in a reactor constructed from natural uranium and a suitable moderator, such as heavy water or graphite, alone.

Uranium-235 can be concentrated by gaseous diffusion and other physical processes, if desired, and used directly as a nuclear fuel, instead of natural uranium, or used as an explosive. Natural uranium, slightly enriched with 235U by a small percentage, is used to fuel nuclear power reactors to generate electricity. Natural thorium can be irradiated with neutrons to produce the important isotope 233U as follows: 232Th(n, gamma) --> 233Th(beta) --> 233Pa(beta) --> 233U. While thorium itself is not fissionable, 233U is, and in this way may be used as a nuclear fuel. One pound of completely fissioned uranium has the fuel value of over 1500 tons of coal. The uses of nuclear fuels to generate electrical power, to make isotopes for peaceful purposes, and to make explosives are well known. Uranium in the U.S.A. is controlled by the U.S. Nuclear Regulatory Commission. New uses are being found for depleted uranium, i.e., uranium with the percentage of 235U lowered to about 0.2%.

Uranium is used in inertial guidance devices, in gyro compasses, as counterweights for aircraft control surfaces, as ballast for missile reentry vehicles, and as a shielding material. Uranium metal is used for X-ray targets for production of high-energy X-rays; the nitrate was once used as a photographic toner, and the acetate was once used in analytical chemistry. Crystals of uranium nitrate are triboluminescent. Uranium salts have also been used for producing yellow "Vaseline" glass and glazes. Hazards Uranium and its compounds are highly toxic, both from a chemical and radiological standpoint. Finely divided uranium metal, being pyrophoric, presents a fire hazard. In nature, U(VI) forms highly soluble carbonate complexes at alkaline pH. This leads to an increase in mobility and availability of uranium to groundwater and soil from nuclear waste repositories which leads to health hazards. Working with uranium requires the knowledge of the maximum allowable concentrations that may be inhaled or ingested. Recently, the natural presence of uranium in many soils has become of concern to homeowners because of the generation of radioactive radon gas and its daughters particularly in confined spaces with low circulation such as basements. Updated by Dr.

David Hobart, July 23, 2013

Question: What element comes between Uranium and Plutonium in the periodic table?
Answer: Neptunium