Organosilicon

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A carbon–silicon bond present in all organosilicon compounds

 

Organosilicon compounds are organometallic compounds containing carbon–silicon bonds. Organosilicon chemistry is the corresponding science of their preparation and properties. Most organosilicon compounds are similar to the ordinary organic compounds, being colourless, flammable, hydrophobic, and stable to air. Silicon carbide is an inorganic compound.

Occurrence and applications

Organosilicon compounds are widely encountered in commercial products. Most common are sealants, caulks, adhesives, and coatings made from silicones. Others important uses include synthesis of polyhedral oligomeric silsesquioxanes, agricultural and plant control adjuvants commonly used in conjunction with herbicides and fungicides.

Silicone caulk, commercial sealants, are mainly composed of organosilicon compounds.

 

Polydimethylsiloxane (PDMS) is the principal component of silicones.

Biology and medicine

Carbon–silicon bonds are absent in biology. Silicates, on the other hand, have known existence in diatoms. Silafluofen is an organosilicon compound that functions as a pyrethroid insecticide. Several organosilicon compounds have been investigated as pharmaceuticals.

Properties of Si–C, Si–O, and Si–F bonds

In most organosilicon compounds, Si is tetravalent with tetrahedral molecular geometry. Carbon–silicon bonds compared to carbon–carbon bonds are longer (186 pm vs. 154 pm) and weaker with bond dissociation energy 451 kJ/mol vs. 607 kJ/mol. The C–Si bond is somewhat polarised towards carbon due to carbon's greater electronegativity (C 2.55 vs Si 1.90). The Si–C bond can be broken more readily than typical C–C bonds. One manifestation of bond polarization in organosilanes is found in the Sakurai reaction. Certain alkyl silanes can be oxidized to an alcohol in the Fleming–Tamao oxidation. 

Another manifestation is the β-silicon effect describes the stabilizing effect of a β-silicon atom on a carbocation with many implications for reactivity. 

Si–O bonds are much stronger (809 kJ/mol compared to 538 kJ/mol) than a typical C–O single bond. The favorable formation of Si–O bonds drives many organic reactions such as the Brook rearrangement and Peterson olefination. Compared to the strong Si–O bond, the Si–F bond is even stronger.

Preparation

The first organosilicon compound, tetraethylsilane, was prepared by Charles Friedel and James Crafts in 1863 by reaction of tetrachlorosilane with diethylzinc. 

The bulk of organosilicon compounds derive from organosilicon chlorides (CH₃)_4-xSiCl_x. These chlorides are produced by the "Direct process", which entails the reaction of methyl chloride with a silicon-copper alloy. The main and most sought-after product is dimethyldichlorosilane:

2 CH₃Cl + Si → (CH₃)₂SiCl₂

A variety of other products are obtained, including trimethylsilyl chloride and methyltrichlorosilane. About 1 million tons of organosilicon compounds are prepared annually by this route. The method can also be used for phenyl chlorosilanes.

Hydrosilylation

After the Direct Process, a second major method for the formation of Si-C bonds is hydrosilylation (also called hydrosilation). In this process, compounds with Si-H bonds (hydrosilanes) add to unsaturated substrates. Commercially, the main substrates are alkenes. Other unsaturated functional groups—alkynes, imines, ketones, and aldehydes. An example is the hydrosilation of phenylacetylene:

Hydrosilylation requires metal catalysts, especially those based on platinum group metals.

In the related silylmetalation, a metal replaces the hydrogen atom.

Functional groups

Silicon is a component of many functional groups. Most of these are analogous to organic compounds. The overarching exception is the rarity of multiple bonds to silicon, as reflected in the double bond rule.

Silanols, siloxides, and siloxanes

Silanols are analogues of alcohols. They are generally prepared by hydrolysis of silyl chlorides:

R₃SiCl + H₂O → R₃SiOH + HCl

Less frequently silanols are prepared by oxidation of silyl hydrides, a reaction that uses a metal catalyst:

2 R₃SiH + O₂ → 2 R₃SiOH

Many silanols have been isolated including (CH₃)₃SiOH and (C₆H₅)₃SiOH. They are about 500x more acidic than the corresponding alcohols. Siloxides are the deprotonated derivatives of silanols:

R₃SiOH + NaOH → R₃SiONa + H₂O

Silanols tend to dehydrate to give siloxanes:

2 R₃SiOH → R₃Si-O-SiR₃ + H₂O

Polymers with repeating siloxane linkages are called silicones. Compounds with an Si=O double bond called silanones are extremely unstable.

Silyl ethers

Silyl ethers have the connectivity Si-O-C. They are typically prepared by the reaction of alcohols with silyl chlorides:

(CH₃)₃SiCl + ROH → (CH₃)₃Si-O-R + HCl

Silyl ethers are extensively used as protective groups for alcohols. 

Exploiting the strength of the Si-F bond, fluoride sources such as tetra-n-butylammonium fluoride (TBAF) are used in deprotection of silyl ethers:

(CH₃)₃Si-O-R + F⁻ + H₂O → (CH₃)₃Si-F + H-O-R + OH⁻

Silyl chlorides

Organosilyl chlorides are important commodity chemicals. They are mainly used to produce silicone polymers as described above. Especially important silyl chlorides are dimethyldichlorosilane (Me₂SiCl₂), methyltrichlorosilane (MeSiCl₃), and trimethylsilyl chloride (Me₃SiCl). More specialized derivatives that find commercial applications include dichloromethylphenylsilane, trichloro(chloromethyl)silane, trichloro(dichlorophenyl)silane, trichloroethylsilane, and phenyltrichlorosilane. 

Although proportionately a minor outlet, organosilicon compounds are widely used in organic synthesis. Notably trimethylsilyl chloride Me₃SiCl is the main silylating agent. One classic method called the Flood reaction for the synthesis of this compound class is by heating hexaalkyldisiloxanes R₃SiOSiR₃ with concentrated sulfuric acid and a sodium halide.

Silyl hydrides

The silicon to hydrogen bond is longer than the C–H bond (148 compared to 105 pm) and weaker (299 compared to 338 kJ/mol). Hydrogen is more electronegative than silicon hence the naming convention of silyl hydrides. Commonly the presence of the hydride is not mentioned in the name of the compound. Triethylsilane has the formula Et₃SiH. Phenylsilane is PhSiH₃. The parent compound SiH₄ is called silane.

Silenes

Organosilicon compounds, unlike their carbon counterparts, do not have a rich double bond chemistry. Compounds with silene Si=C bonds (also known as alkylidenesilanes) are laboratory curiosities such as the silicon benzene analogue silabenzene. In 1967, Gusel'nikov and Flowers provided the first evidence for silenes from pyrolysis of dimethylsilacyclobutane. The first stable (kinetically shielded) silene was reported in 1981 by Brook.

Disilenes have Si=Si double bonds and disilynes are silicon analogues of an alkyne. The first Silyne (with a silicon to carbon triple bond) was reported in 2010.

Siloles

Chemical structure of silole

 

Siloles, also called silacyclopentadienes, are members of a larger class of compounds called metalloles. They are the silicon analogs of cyclopentadienes and are of current academic interest due to their electroluminescence and other electronic properties. Siloles are efficient in electron transport. They owe their low lying LUMO to a favorable interaction between the antibonding sigma silicon orbital with an antibonding pi orbital of the butadiene fragment.

Hypercoordinated silicon

Unlike carbon, silicon compounds can be coordinated to five atoms as well in a group of compounds ranging from so-called silatranes, such as phenylsilatrane, to a uniquely stable pentaorganosilicate:

The stability of hypervalent silicon is the basis of the Hiyama coupling, a coupling reaction used in certain specialized organic synthetic applications. The reaction begins with the activation of Si-C bond by fluoride:

R-SiR'₃ + R"-X + F⁻ → R-R" + R'₃SiF + X⁻

Various reactions

Certain allyl silanes can be prepared from allylic esters such as 1 and monosilylcopper compounds such as 2 in.

In this reaction type silicon polarity is reversed in a chemical bond with zinc and a formal allylic substitution on the benzoyloxy group takes place.

Environmental effects

Organosilicon compounds affect bee (and other insect) immune expression, making them more susceptible to viral infection.

Copernican principle

From Wikipedia, the free encyclopedia

Figure 'M' (for Latin Mundus) from Johannes Kepler's 1617–1621 Epitome Astronomiae Copernicanae, showing the Earth as belonging to just one of any number of similar stars.

 

In physical cosmology, the Copernican principle states that humans, on the Earth or in the Solar System, are not privileged observers of the universe.

Named for Copernican heliocentrism, it is a working assumption that arises from a modified cosmological extension of Copernicus's argument of a moving Earth. In some sense, it is equivalent to the mediocrity principle.

Origin and implications

Hermann Bondi named the principle after Copernicus in the mid-20th century, although the principle itself dates back to the 16th-17th century paradigm shift away from the Ptolemaic system, which placed Earth at the center of the universe. Copernicus proposed that the motion of the planets can be explained by reference to an assumption that the Sun and not Earth is centrally located and stationary. He argued that the apparent retrograde motion of the planets is an illusion caused by Earth's movement around the Sun, which the Copernican model placed at the centre of the universe. Copernicus himself was mainly motivated by technical dissatisfaction with the earlier system and not by support for any mediocrity principle. In fact, although the Copernican heliocentric model is often described as "demoting" Earth from its central role it had in the Ptolemaic geocentric model, it was successors to Copernicus, notably the 16th century Giordano Bruno, who adopted this new perspective. The Earth's central position had been interpreted as being in the "lowest and filthiest parts". Instead, as Galileo said, the Earth is part of the "dance of the stars" rather than the "sump where the universe's filth and ephemera collect". In the late 20th Century, Carl Sagan asked, "Who are we? We find that we live on an insignificant planet of a humdrum star lost in a galaxy tucked away in some forgotten corner of a universe in which there are far more galaxies than people."

In cosmology, if one assumes the Copernican principle and observes that the universe appears isotropic or the same in all directions from the vantage point of Earth, then one can infer that the universe is generally homogeneous or the same everywhere (at any given time) and is also isotropic about any given point. These two conditions make up the cosmological principle. In practice, astronomers observe that the universe has heterogeneous or non-uniform structures up to the scale of galactic superclusters, filaments and great voids. It becomes more and more homogeneous and isotropic when observed on larger and larger scales, with little detectable structure on scales of more than about 200 million parsecs. However, on scales comparable to the radius of the observable universe, we see systematic changes with distance from Earth. For instance, galaxies contain more young stars and are less clustered, and quasars appear more numerous. While this might suggest that Earth is at the center of the universe, the Copernican principle requires us to interpret it as evidence for the evolution of the universe with time: this distant light has taken most of the age of the universe to reach Earth and shows the universe when it was young. The most distant light of all, cosmic microwave background radiation, is isotropic to at least one part in a thousand.

Modern mathematical cosmology is based on the assumption that the Cosmological principle is almost, but not exactly, true on the largest scales. The Copernican principle represents the irreducible philosophical assumption needed to justify this, when combined with the observations. 

Michael Rowan-Robinson emphasizes the Copernican principle as the threshold test for modern thought, asserting that: "It is evident that in the post-Copernican era of human history, no well-informed and rational person can imagine that the Earth occupies a unique position in the universe."

Bondi and Thomas Gold used the Copernican principle to argue for the perfect cosmological principle which maintains that the universe is also homogeneous in time, and is the basis for the steady-state cosmology. However, this strongly conflicts with the evidence for cosmological evolution mentioned earlier: the universe has progressed from extremely different conditions at the Big Bang, and will continue to progress toward extremely different conditions, particularly under the rising influence of dark energy, apparently toward the Big Freeze or Big Rip. 

Since the 1990s the term has been used (interchangeably with "the Copernicus method") for J. Richard Gott's Bayesian-inference-based prediction of duration of ongoing events, a generalized version of the Doomsday argument.

Derived equations

Copernicus was able to derive the lengths of superior and inferior planet's orbits, based on the lengths of their sidereal orbits.

Where for superior planets,

${\displaystyle {\frac {1}{P}}+{\frac {1}{S}}={\frac {1}{E}}}$

with P is the sidereal period of the superior planet, S the synodic period of the planet, and E the period of the planet that measurements are based on, in most cases the Earth.

The Formula for inferior planets is 

${\displaystyle {\frac {1}{P}}-{\frac {1}{S}}={\frac {1}{E}}}$

Tests of the principle

The Copernican principle has never been proven, and in the most general sense cannot be proven, but it is implicit in many modern theories of physics. Cosmological models are often derived with reference to the Cosmological principle, slightly more general than the Copernican principle, and many tests of these models can be considered tests of the Copernican principle.

Historical

Before the term Copernican principle was even coined, Earth was repeatedly shown not to have any special location in the universe. The Copernican Revolution dethroned Earth to just one of many planets orbiting the Sun. Proper motion was mentioned by Halley. William Herschel found that the Solar System is moving through space within our disk-shaped Milky Way galaxy. Edwin Hubble showed that the Milky Way galaxy is just one of many galaxies in the universe. Examination of the galaxy's position and motion in the universe led to the Big Bang theory and the whole of modern cosmology.

Modern tests

Recent and planned tests relevant to the cosmological and Copernican principles include

• time drift of cosmological redshifts;
• modelling the local gravitational potential using reflection of cosmic microwave background (CMB) photons;
• the redshift dependence of the luminosity of supernovae;
• the kinetic Sunyaev-Zel’dovich effect in relation to dark energy;
• cosmic neutrino background;
• the integrated Sachs-Wolfe effect;
• testing the isotropy and homogeneity of the CMB.

Physics without the principle

The standard model of cosmology, the Lambda-CDM model, assumes the Copernican principle and the more general Cosmological principle and observations are largely consistent but there are always unsolved problems. Some cosmologists and theoretical physicists design models lacking the Cosmological or Copernican principles, to constrain the valid values of observational results, to address specific known issues, and to propose tests to distinguish between current models and other possible models. 

A prominent example in this context is the observed accelerating universe and the cosmological constant issue. An alternative proposal to dark energy is that the universe is much more inhomogeneous than currently assumed, and specifically that we are in an extremely large low-density void. To match observations we would have to be very close to the centre of this void, immediately contradicting the Copernican principle.

Actinide

From Wikipedia, the free encyclopedia

The atomic bomb dropped on Nagasaki had a plutonium charge.

 

The actinide /ˈæktɪnaɪd/ or actinoid /ˈæktɪnɔɪd/ (IUPAC nomenclature) series encompasses the 15 metallic chemical elements with atomic numbers from 89 to 103, actinium through lawrencium.

Strictly speaking, both actinium and lawrencium have been labeled as group 3 elements, but both elements are often included in any general discussion of the chemistry of the actinide elements. Actinium is the more often omitted of the two, because its placement as a group 3 element is somewhat more common in texts and for semantic reasons: since "actinide" means "like actinium", it has been argued that actinium cannot logically be an actinide, even though IUPAC acknowledges its inclusion based on common usage.

The actinide series derives its name from the first element in the series, actinium. The informal chemical symbol An is used in general discussions of actinide chemistry to refer to any actinide. All but one of the actinides are f-block elements, with the exception being either actinium or lawrencium. The series mostly corresponds to the filling of the 5f electron shell, although actinium and thorium lack any f-electrons, and curium and lawrencium have the same number as the preceding element. In comparison with the lanthanides, also mostly f-block elements, the actinides show much more variable valence. They all have very large atomic and ionic radii and exhibit an unusually large range of physical properties. While actinium and the late actinides (from americium onwards) behave similarly to the lanthanides, the elements thorium, protactinium, and uranium are much more similar to transition metals in their chemistry, with neptunium and plutonium occupying an intermediate position.

All actinides are radioactive and release energy upon radioactive decay; naturally occurring uranium and thorium, and synthetically produced plutonium are the most abundant actinides on Earth. These are used in nuclear reactors and nuclear weapons. Uranium and thorium also have diverse current or historical uses, and americium is used in the ionization chambers of most modern smoke detectors. 

Of the actinides, primordial thorium and uranium occur naturally in substantial quantities. The radioactive decay of uranium produces transient amounts of actinium and protactinium, and atoms of neptunium and plutonium are occasionally produced from transmutation reactions in uranium ores. The other actinides are purely synthetic elements. Nuclear weapons tests have released at least six actinides heavier than plutonium into the environment; analysis of debris from a 1952 hydrogen bomb explosion showed the presence of americium, curium, berkelium, californium, einsteinium and fermium.

In presentations of the periodic table, the lanthanides and the actinides are customarily shown as two additional rows below the main body of the table, with placeholders or else a selected single element of each series (either lanthanum or lutetium, and either actinium or lawrencium, respectively) shown in a single cell of the main table, between barium and hafnium, and radium and rutherfordium, respectively. This convention is entirely a matter of aesthetics and formatting practicality; a rarely used wide-formatted periodic table inserts the lanthanide and actinide series in their proper places, as parts of the table's sixth and seventh rows (periods).

Discovery, isolation and synthesis

Synthesis of transuranium elements

Element	Year	Method
Neptunium	1940	Bombarding 238U by neutrons
Plutonium	1941	Bombarding 238U by deuterons
Americium	1944	Bombarding 239Pu by neutrons
Curium	1944	Bombarding 239Pu by α-particles
Berkelium	1949	Bombarding 241Am by α-particles
Californium	1950	Bombarding 242Cm by α-particles
Einsteinium	1952	As a product of nuclear explosion
Fermium	1952	As a product of nuclear explosion
Mendelevium	1955	Bombarding 253Es by α-particles
Nobelium	1965	Bombarding 243Am by 15N or 238U with 22Ne
Lawrencium	1961 –1971	Bombarding 252Cf by 10B or 11B and of 243Am with 18O

Like the lanthanides, the actinides form a family of elements with similar properties. Within the actinides, there are two overlapping groups: transuranium elements, which follow uranium in the periodic table—and transplutonium elements, which follow plutonium. Compared to the lanthanides, which (except for promethium) are found in nature in appreciable quantities, most actinides are rare. The majority of them do not even occur in nature, and of those that do, only thorium and uranium do so in more than trace quantities. The most abundant or easily synthesized actinides are uranium and thorium, followed by plutonium, americium, actinium, protactinium, neptunium, and curium.^\

 

The existence of transuranium elements was suggested by Enrico Fermi based on his experiments in 1934. However, even though four actinides were known by that time, it was not yet understood that they formed a family similar to lanthanides. The prevailing view that dominated early research into transuranics was that they were regular elements in the 7th period, with thorium, protactinium and uranium corresponding to 6th-period hafnium, tantalum and tungsten, respectively. Synthesis of transuranics gradually undermined this point of view. By 1944 an observation that curium failed to exhibit oxidation states above 4 (whereas its supposed 6th period homolog, platinum, can reach oxidation state of 6) prompted Glenn Seaborg to formulate a so-called "actinide hypothesis". Studies of known actinides and discoveries of further transuranic elements provided more data in support of this point of view, but the phrase "actinide hypothesis" (the implication being that a "hypothesis" is something that has not been decisively proven) remained in active use by scientists through the late 1950s.

At present, there are two major methods of producing isotopes of transplutonium elements: (1) irradiation of the lighter elements with either neutrons or (2) accelerated charged particles. The first method is most important for applications, as only neutron irradiation using nuclear reactors allows the production of sizeable amounts of synthetic actinides; however, it is limited to relatively light elements. The advantage of the second method is that elements heavier than plutonium, as well as neutron-deficient isotopes, can be obtained, which are not formed during neutron irradiation.

In 1962–1966, there were attempts in the United States to produce transplutonium isotopes using a series of six underground nuclear explosions. Small samples of rock were extracted from the blast area immediately after the test to study the explosion products, but no isotopes with mass number greater than 257 could be detected, despite predictions that such isotopes would have relatively long half-lives of α-decay. This non-observation was attributed to spontaneous fission owing to the large speed of the products and to other decay channels, such as neutron emission and nuclear fission.

From actinium to uranium

Enrico Fermi suggested the existence of transuranium elements in 1934.

 

Uranium and thorium were the first actinides discovered. Uranium was identified in 1789 by the German chemist Martin Heinrich Klaproth in pitchblende ore. He named it after the planet Uranus, which had been discovered only eight years earlier. Klaproth was able to precipitate a yellow compound (likely sodium diuranate) by dissolving pitchblende in nitric acid and neutralizing the solution with sodium hydroxide. He then reduced the obtained yellow powder with charcoal, and extracted a black substance that he mistook for metal. Only 60 years later, the French scientist Eugène-Melchior Péligot identified it as uranium oxide. He also isolated the first sample of uranium metal by heating uranium tetrachloride with metallic potassium. The atomic mass of uranium was then calculated as 120, but Dmitri Mendeleev in 1872 corrected it to 240 using his periodicity laws. This value was confirmed experimentally in 1882 by K. Zimmerman.

Thorium oxide was discovered by Friedrich Wöhler in the mineral Thorianite, which was found in Norway (1827). Jöns Jacob Berzelius characterized this material in more detail by in 1828. By reduction of thorium tetrachloride with potassium, he isolated the metal and named it thorium after the Norse god of thunder and lightning Thor. The same isolation method was later used by Péligot for uranium.

Actinium was discovered in 1899 by André-Louis Debierne, an assistant of Marie Curie, in the pitchblende waste left after removal of radium and polonium. He described the substance (in 1899) as similar to titanium and (in 1900) as similar to thorium. The discovery of actinium by Debierne was however questioned in 1971 and 2000, arguing that Debierne's publications in 1904 contradicted his earlier work of 1899–1900. This view instead credits the 1902 work of Friedrich Oskar Giesel, who discovered a radioactive element named emanium that behaved similarly to lanthanum. The name actinium comes from the Greek aktis, aktinos (ακτίς, ακτίνος), meaning beam or ray. This metal was discovered not by its own radiation but by the radiation of the daughter products. Owing to the close similarity of actinium and lanthanum and low abundance, pure actinium could only be produced in 1950. The term actinide was probably introduced by Victor Goldschmidt in 1937.

Protactinium was possibly isolated in 1900 by William Crookes. It was first identified in 1913, when Kasimir Fajans and Oswald Helmuth Göhring encountered the short-lived isotope ^234mPa (half-life 1.17 minutes) during their studies of the ²³⁸U decay. They named the new element brevium (from Latin brevis meaning brief); the name was changed to protoactinium (from Greek πρῶτος + ἀκτίς meaning "first beam element") in 1918 when two groups of scientists, led by the Austrian Lise Meitner and Otto Hahn of Germany and Frederick Soddy and John Cranston of Great Britain, independently discovered the much longer-lived ²³¹Pa. The name was shortened to protactinium in 1949. This element was little characterized until 1960, when A. G. Maddock and his co-workers in the U.K. isolated 130 grams of protactinium from 60 tonnes of waste left after extraction of uranium from its ore.

Neptunium and above

Neptunium (named for the planet Neptune, the next planet out from Uranus, after which uranium was named) was discovered by Edwin McMillan and Philip H. Abelson in 1940 in Berkeley, California. They produced the ²³⁹Np isotope (half-life = 2.4 days) by bombarding uranium with slow neutrons. It was the first transuranium element produced synthetically.

Glenn T. Seaborg and his group at the University of California at Berkeley synthesized Pu, Am, Cm, Bk, Cf, Es, Fm, Md, No and element 106, which was later named seaborgium in his honor while he was still living. They also synthesized more than a hundred actinide isotopes.

 

Transuranium elements do not occur in sizeable quantities in nature and are commonly synthesized via nuclear reactions conducted with nuclear reactors. For example, under irradiation with reactor neutrons, uranium-238 partially converts to plutonium-239:

${\displaystyle {\ce {{^{238}_{92}U}+{}_{0}^{1}n->{}_{92}^{239}U->[\beta ^{-}][23.5\ {\ce {min}}]{}_{93}^{239}Np->[\beta ^{-}][2.3\ {\ce {days}}]{}_{94}^{239}Pu}}\left({\ce {->[\alpha ][2.4\cdot 10^{4}\ {\ce {years}}]}}\right){\ce {^{235}_{92}U}}}$

This synthesis reaction was used by Fermi and his collaborators in their design of the reactors located at the Hanford Site, which produced significant amounts of plutonium-239 for the nuclear weapons of the Manhattan Project and the United States' post-war nuclear arsenal.

Actinides with the highest mass numbers are synthesized by bombarding uranium, plutonium, curium and californium with ions of nitrogen, oxygen, carbon, neon or boron in a particle accelerator. So, nobelium was produced by bombarding uranium-238 with neon-22 as

${\displaystyle {\ce {_{92}^{238}U + _{10}^{22}Ne -> _{102}^{256}No + 4_0^1n}}}$.

The first isotopes of transplutonium elements, americium-241 and curium-242, were synthesized in 1944 by Glenn T. Seaborg, Ralph A. James and Albert Ghiorso. Curium-242 was obtained by bombarding plutonium-239 with 32-MeV α-particles

${\displaystyle {\ce {_{94}^{239}Pu + _2^4He -> _{96}^{242}Cm + _0^1n}}}$.

The americium-241 and curium-242 isotopes also were produced by irradiating plutonium in a nuclear reactor. The latter element was named after Marie Curie and her husband Pierre who are noted for discovering radium and for their work in radioactivity.

Bombarding curium-242 with α-particles resulted in an isotope of californium ²⁴⁵Cf (1950), and a similar procedure yielded in 1949 berkelium-243 from americium-241. The new elements were named after Berkeley, California, by analogy with its lanthanide homologue terbium, which was named after the village of Ytterby in Sweden.

In 1945, B. B. Cunningham obtained the first bulk chemical compound of a transplutonium element, namely americium hydroxide. Over the next three to four years, milligram quantities of americium and microgram amounts of curium were accumulated that allowed production of isotopes of berkelium (Thomson, 1949) and californium (Thomson, 1950). Sizeable amounts of these elements were produced only in 1958 (Burris B. Cunningham and Stanley G. Thomson), and the first californium compound (0.3 µg of CfOCl) was obtained only in 1960 by B. B. Cunningham and J. C. Wallmann.

Einsteinium and fermium were identified in 1952–1953 in the fallout from the "Ivy Mike" nuclear test (1 November 1952), the first successful test of a hydrogen bomb. Instantaneous exposure of uranium-238 to a large neutron flux resulting from the explosion produced heavy isotopes of uranium, including uranium-253 and uranium-255, and their β-decay yielded einsteinium-253 and fermium-255. The discovery of the new elements and the new data on neutron capture were initially kept secret on the orders of the U.S. military until 1955 due to Cold War tensions. Nevertheless, the Berkeley team were able to prepare einsteinium and fermium by civilian means, through the neutron bombardment of plutonium-239, and published this work in 1954 with the disclaimer that it was not the first studies that had been carried out on the elements. The "Ivy Mike" studies were declassified and published in 1955. The first significant (submicrograms) amounts of einsteinium were produced in 1961 by Cunningham and colleagues, but this has not been done for fermium yet.

The first isotope of mendelevium, ²⁵⁶Md (half-life 87 min), was synthesized by Albert Ghiorso, Glenn T. Seaborg, Gregory R. Choppin, Bernard G. Harvey and Stanley G. Thompson when they bombarded an ²⁵³Es target with alpha particles in the 60-inch cyclotron of Berkeley Radiation Laboratory; this was the first isotope of any element to be synthesized one atom at a time.

There were several attempts to obtain isotopes of nobelium by Swedish (1957) and American (1958) groups, but the first reliable result was the synthesis of ²⁵⁶No by the Russian group (Georgy Flyorov et al.) in 1965, as acknowledged by the IUPAC in 1992. In their experiments, Flyorov et al. bombarded uranium-238 with neon-22.

In 1961, Ghiorso et al. obtained the first isotope of lawrencium by irradiating californium (mostly californium-252) with boron-10 and boron-11 ions. The mass number of this isotope was not clearly established (possibly 258 or 259) at the time. In 1965, ²⁵⁶Lr was synthesized by Flyorov et al. from ²⁴³Am and ¹⁸O. Thus IUPAC recognized the nuclear physics teams at Dubna and Berkeley as the co-discoverers of lawrencium.

Isotopes

Actinides have 89−103 protons and usually 117−159 neutrons.

 

Thirty-one isotopes of actinium and eight excited isomeric states of some of its nuclides were identified by 2010. Three isotopes, ²²⁵Ac, ²²⁷Ac and ²²⁸Ac, were found in nature and the others were produced in the laboratory; only the three natural isotopes are used in applications. Actinium-225 is a member of the radioactive neptunium series; it was first discovered in 1947 as a decay product of uranium-233, it is an α-emitter with a half-life of 10 days. Actinium-225 is less available than actinium-228, but is more promising in radiotracer applications. Actinium-227 (half-life 21.77 years) occurs in all uranium ores, but in small quantities. One gram of uranium (in radioactive equilibrium) contains only 2×10⁻¹⁰ gram of ²²⁷Ac. Actinium-228 is a member of the radioactive thorium series formed by the decay of ²²⁸Ra; it is a β⁻ emitter with a half-life of 6.15 hours. In one tonne of thorium there is 5×10⁻⁸ gram of ²²⁸Ac. It was discovered by Otto Hahn in 1906.

28 isotopes of protactinium are known with mass numbers 212–239 as well as three excited isomeric states. Only ²³¹Pa and ²³⁴Pa have been found in nature. All the isotopes have short lifetime, except for protactinium-231 (half-life 32,760 years). The most important isotopes are ²³¹Pa and ²³³Pa, which is an intermediate product in obtaining uranium-233 and is the most affordable among artificial isotopes of protactinium. ²³³Pa has convenient half-life and energy of γ-radiation, and thus was used in most studies of protactinium chemistry. Protactinium-233 is a β-emitter with a half-life of 26.97 days.

Uranium has the highest number (25) of both natural and synthetic isotopes. They have mass numbers of 215–242 (except 220 and 241), and three of them, ²³⁴U, ²³⁵U and ²³⁸U, are present in appreciable quantities in nature. Among others, the most important is ²³³U, which is a final product of transformations of ²³²Th irradiated by slow neutrons. ²³³U has a much higher fission efficiency by low-energy (thermal) neutrons, compared e.g. with ²³⁵U. Most uranium chemistry studies were carried out on uranium-238 owing to its long half-life of 4.4×10⁹ years.

There are 23 isotopes of neptunium with mass numbers of 219 and 223–244; they are all highly radioactive. The most popular among scientists are long-lived ²³⁷Np (t_1/2 = 2.20×10⁶ years) and short-lived ²³⁹Np, ²³⁸Np (t_1/2 ~ 2 days).

Eighteen isotopes of americium are known with mass numbers from 229 to 247 (with the exception of 231). The most important are ²⁴¹Am and ²⁴³Am, which are alpha-emitters and also emit soft, but intense γ-rays; both of them can be obtained in an isotopically pure form. Chemical properties of americium were first studied with ²⁴¹Am, but later shifted to ²⁴³Am, which is almost 20 times less radioactive. The disadvantage of ²⁴³Am is production of the short-lived daughter isotope ²³⁹Np, which has to be considered in the data analysis.

Among 19 isotopes of curium, the most accessible are ²⁴²Cm and ²⁴⁴Cm; they are α-emitters, but with much shorter lifetime than the americium isotopes. These isotopes emit almost no γ-radiation, but undergo spontaneous fission with the associated emission of neutrons. More long-lived isotopes of curium (^245–248Cm, all α-emitters) are formed as a mixture during neutron irradiation of plutonium or americium. Upon short irradiation, this mixture is dominated by ²⁴⁶Cm, and then ²⁴⁸Cm begins to accumulate. Both of these isotopes, especially ²⁴⁸Cm, have a longer half-life (3.48×10⁵ years) and are much more convenient for carrying out chemical research than ²⁴²Cm and ²⁴⁴Cm, but they also have a rather high rate of spontaneous fission. ²⁴⁷Cm has the longest lifetime among isotopes of curium (1.56×10⁷ years), but is not formed in large quantities because of the strong fission induced by thermal neutrons. 

Eighteen isotopes of berkelium were identified with mass numbers 233–234, 236, and 238–252. Only ²⁴⁹Bk is available in large quantities; it has a relatively short half-life of 330 days and emits mostly soft β-particles, which are inconvenient for detection. Its alpha radiation is rather weak (1.45×10⁻³% with respect to β-radiation), but is sometimes used to detect this isotope. ²⁴⁷Bk is an alpha-emitter with a long half-life of 1,380 years, but it is hard to obtain in appreciable quantities; it is not formed upon neutron irradiation of plutonium because of the β-stability of isotopes of curium isotopes with mass number below 248.

Isotopes of californium with mass numbers 237–256 are formed in nuclear reactors; californium-253 is a β-emitter and the rest are α-emitters. The isotopes with even mass numbers (²⁵⁰Cf, ²⁵²Cf and ²⁵⁴Cf) have a high rate of spontaneous fission, especially ²⁵⁴Cf of which 99.7% decays by spontaneous fission. Californium-249 has a relatively long half-life (352 years), weak spontaneous fission and strong γ-emission that facilitates its identification. ²⁴⁹Cf is not formed in large quantities in a nuclear reactor because of the slow β-decay of the parent isotope ²⁴⁹Bk and a large cross section of interaction with neutrons, but it can be accumulated in the isotopically pure form as the β-decay product of (pre-selected) ²⁴⁹Bk. Californium produced by reactor-irradiation of plutonium mostly consists of ²⁵⁰Cf and ²⁵²Cf, the latter being predominant for large neutron fluences, and its study is hindered by the strong neutron radiation.

Properties of some transplutonium isotope pairs

Parent isotope	t1/2	Daughter isotope	t1/2	Time to establish radioactive equilibrium
243Am	7370 years	239Np	2.35 days	47.3 days
245Cm	8265 years	241Pu	14 years	129 years
247Cm	1.64×107 years	243Pu	4.95 hours	7.2 days
254Es	270 days	250Bk	3.2 hours	35.2 hours
255Es	39.8 days	255Fm	22 hours	5 days
257Fm	79 days	253Cf	17.6 days	49 days

Among the 18 known isotopes of einsteinium with mass numbers from 240 to 257, the most affordable is ²⁵³Es. It is an α-emitter with a half-life of 20.47 days, a relatively weak γ-emission and small spontaneous fission rate as compared with the isotopes of californium. Prolonged neutron irradiation also produces a long-lived isotope ²⁵⁴Es (t_1/2 = 275.5 days).

Twenty isotopes of fermium are known with mass numbers of 241–260. ²⁵⁴Fm, ²⁵⁵Fm and ²⁵⁶Fm are α-emitters with a short half-life (hours), which can be isolated in significant amounts. ²⁵⁷Fm (t_1/2 = 100 days) can accumulate upon prolonged and strong irradiation. All these isotopes are characterized by high rates of spontaneous fission.

Among the 16 known isotopes of mendelevium (mass numbers from 245 to 260), the most studied is ²⁵⁶Md, which mainly decays through the electron capture (α-radiation is ≈10%) with the half-life of 77 minutes. Another alpha emitter, ²⁵⁸Md, has a half-life of 53 days. Both these isotopes are produced from rare einsteinium (²⁵³Es and ²⁵⁵Es respectively), that therefore limits their availability.

Long-lived isotopes of nobelium and isotopes of lawrencium (and of heavier elements) have relatively short half-lives. For nobelium, 11 isotopes are known with mass numbers 250–260 and 262. The chemical properties of nobelium and lawrencium were studied with ²⁵⁵No (t_1/2 = 3 min) and ²⁵⁶Lr (t_1/2 = 35 s). The longest-lived nobelium isotope, ²⁵⁹No, has a half-life of approximately 1 hour.

Among all of these, the only isotopes that occur in sufficient quantities in nature to be detected in anything more than traces and have a measurable contribution to the atomic weights of the actinides are the primordial ²³²Th, ²³⁵U, and ²³⁸U, and three long-lived decay products of natural uranium, ²³⁰Th, ²³¹Pa, and ²³⁴U. Natural thorium consists of 0.02(2)% ²³⁰Th and 99.98(2)% ²³²Th; natural protactinium consists of 100% ²³¹Pa; and natural uranium consists of 0.0054(5)% ²³⁴U, 0.7204(6)% ²³⁵U, and 99.2742(10)% ²³⁸U.

Distribution in nature

Unprocessed uranium ore

 

Thorium and uranium are the most abundant actinides in nature with the respective mass concentrations of 16 ppm and 4 ppm. Uranium mostly occurs in the Earth's crust as a mixture of its oxides in the minerals uraninite, which is also called pitchblende because of its black color. There are several dozens of other uranium minerals such as carnotite (KUO₂VO₄·3H₂O) and autunite (Ca(UO₂)₂(PO₄)₂·nH₂O). The isotopic composition of natural uranium is ²³⁸U (relative abundance 99.2742%), ²³⁵U (0.7204%) and ²³⁴U (0.0054%); of these ²³⁸U has the largest half-life of 4.51×10⁹ years. The worldwide production of uranium in 2009 amounted to 50,572 tonnes, of which 27.3% was mined in Kazakhstan. Other important uranium mining countries are Canada (20.1%), Australia (15.7%), Namibia (9.1%), Russia (7.0%), and Niger (6.4%).

Content of plutonium in uranium and thorium ores

Ore	Location	Uranium content, %	Mass ratio 239Pu/ore	Ratio 239Pu/U (×1012)
Uraninite	Canada	13.5	9.1×10−12	7.1
Uraninite	Congo	38	4.8×10−12	12
Uraninite	Colorado, US	50	3.8×10−12	7.7
Monazite	Brazil	0.24	2.1×10−14	8.3
Monazite	North Carolina, US	1.64	5.9×10−14	3.6
Fergusonite	-	0.25	<1 span="" style="margin: 0 .15em 0 .25em;">×10−14

<4 span=""> Carnotite - 10 <4 span="" style="margin: 0 .15em 0 .25em;">× 10⁻¹⁴ <0 .4="" span="">

The most abundant thorium minerals are thorianite (ThO₂), thorite (ThSiO₄) and monazite, ((Th,Ca,Ce)PO₄). Most thorium minerals contain uranium and vice versa; and they all have significant fraction of lanthanides. Rich deposits of thorium minerals are located in the United States (440,000 tonnes), Australia and India (~300,000 tonnes each) and Canada (~100,000 tonnes).

The abundance of actinium in the Earth's crust is only about 5×10⁻¹⁵%. Actinium is mostly present in uranium-containing, but also in other minerals, though in much smaller quantities. The content of actinium in most natural objects corresponds to the isotopic equilibrium of parent isotope ²³⁵U, and it is not affected by the weak Ac migration. Protactinium is more abundant (10⁻¹²%) in the Earth's crust than actinium. It was discovered in the uranium ore in 1913 by Fajans and Göhring. As actinium, the distribution of protactinium follows that of ²³⁵U.

The half-life of the longest-lived isotope of neptunium, ²³⁷Np, is negligible compared to the age of the Earth. Thus neptunium is present in nature in negligible amounts produced as intermediate decay products of other isotopes. Traces of plutonium in uranium minerals were first found in 1942, and the more systematic results on ²³⁹Pu are summarized in the table (no other plutonium isotopes could be detected in those samples). The upper limit of abundance of the longest-living isotope of plutonium, ²⁴⁴Pu, is 3×10⁻²⁰%. Plutonium could not be detected in samples of lunar soil. Owing to its scarcity in nature, most plutonium is produced synthetically.

Extraction

Monazite: a major thorium mineral

 

Owing to the low abundance of actinides, their extraction is a complex, multistep process. Fluorides of actinides are usually used because they are insoluble in water and can be easily separated with redox reactions. Fluorides are reduced with calcium, magnesium or barium:

${\displaystyle {\begin{array}{l}{}\\{\ce {2AmF3{}+3Ba->[{\ce {1150-1350^{\circ }C}}]3BaF2{}+2Am}}\\{\ce {PuF4{}+2Ba->[{\ce {1200^{\circ }C}}]2BaF2{}+Pu}}\\{\ce {UF4{}+2Mg->[{\ce {>500^{\circ }C}}]U{}+2MgF2}}\\{}\end{array}}}$

Among the actinides, thorium and uranium are the easiest to isolate. Thorium is extracted mostly from monazite: thorium pyrophosphate (ThP₂O₇) is reacted with nitric acid, and the produced thorium nitrate treated with tributyl phosphate. Rare-earth impurities are separated by increasing the pH in sulfate solution.

In another extraction method, monazite is decomposed with a 45% aqueous solution of sodium hydroxide at 140 °C. Mixed metal hydroxides are extracted first, filtered at 80 °C, washed with water and dissolved with concentrated hydrochloric acid. Next, the acidic solution is neutralized with hydroxides to pH = 5.8 that results in precipitation of thorium hydroxide (Th(OH)₄) contaminated with ~3% of rare-earth hydroxides; the rest of rare-earth hydroxides remains in solution. Thorium hydroxide is dissolved in an inorganic acid and then purified from the rare earth elements. An efficient method is the dissolution of thorium hydroxide in nitric acid, because the resulting solution can be purified by extraction with organic solvents:

Separation of uranium and plutonium from nuclear fuel

Th(OH)₄ + 4 HNO₃ → Th(NO₃)₄ + 4 H₂O

Metallic thorium is separated from the anhydrous oxide, chloride or fluoride by reacting it with calcium in an inert atmosphere:

ThO₂ + 2 Ca → 2 CaO + Th

Sometimes thorium is extracted by electrolysis of a fluoride in a mixture of sodium and potassium chloride at 700–800 °C in a graphite crucible. Highly pure thorium can be extracted from its iodide with the crystal bar process.

Uranium is extracted from its ores in various ways. In one method, the ore is burned and then reacted with nitric acid to convert uranium into a dissolved state. Treating the solution with a solution of tributyl phosphate (TBP) in kerosene transforms uranium into an organic form UO₂(NO₃)₂(TBP)₂. The insoluble impurities are filtered and the uranium is extracted by reaction with hydroxides as (NH₄)₂U₂O₇ or with hydrogen peroxide as UO₄·2H₂O.

When the uranium ore is rich in such minerals as dolomite, magnesite, etc., those minerals consume much acid. In this case, the carbonate method is used for uranium extraction. Its main component is an aqueous solution of sodium carbonate, which converts uranium into a complex [UO₂(CO₃)₃]⁴⁻, which is stable in aqueous solutions at low concentrations of hydroxide ions. The advantages of the sodium carbonate method are that the chemicals have low corrosivity (compared to nitrates) and that most non-uranium metals precipitate from the solution. The disadvantage is that tetravalent uranium compounds precipitate as well. Therefore, the uranium ore is treated with sodium carbonate at elevated temperature and under oxygen pressure:

2 UO₂ + O₂ + 6 CO²⁻
₃ → 2 [UO₂(CO₃)₃]⁴⁻

This equation suggests that the best solvent for the uranium carbonate processing is a mixture of carbonate with bicarbonate. At high pH, this results in precipitation of diuranate, which is treated with hydrogen in the presence of nickel yielding an insoluble uranium tetracarbonate.

Another separation method uses polymeric resins as a polyelectrolyte. Ion exchange processes in the resins result in separation of uranium. Uranium from resins is washed with a solution of ammonium nitrate or nitric acid that yields uranyl nitrate, UO₂(NO₃)₂·6H₂O. When heated, it turns into UO₃, which is converted to UO₂ with hydrogen:

UO₃ + H₂ → UO₂ + H₂O

Reacting uranium dioxide with hydrofluoric acid changes it to uranium tetrafluoride, which yields uranium metal upon reaction with magnesium metal:

4 HF + UO₂ → UF₄ + 2 H₂O

To extract plutonium, neutron-irradiated uranium is dissolved in nitric acid, and a reducing agent (FeSO₄, or H₂O₂) is added to the resulting solution. This addition changes the oxidation state of plutonium from +6 to +4, while uranium remains in the form of uranyl nitrate (UO₂(NO₃)₂). The solution is treated with a reducing agent and neutralized with ammonium carbonate to pH = 8 that results in precipitation of Pu⁴⁺ compounds.

In another method, Pu⁴⁺ and UO²⁺
₂ are first extracted with tributyl phosphate, then reacted with hydrazine washing out the recovered plutonium.

The major difficulty in separation of actinium is the similarity of its properties with those of lanthanum. Thus actinium is either synthesized in nuclear reactions from isotopes of radium or separated using ion-exchange procedures.

Properties

Actinides have similar properties to lanthanides. The 6d and 7s electronic shells are filled in actinium and thorium, and the 5f shell is being filled with further increase in atomic number; the 4f shell is filled in the lanthanides. The first experimental evidence for the filling of the 5f shell in actinides was obtained by McMillan and Abelson in 1940. As in lanthanides (see lanthanide contraction), the ionic radius of actinides monotonically decreases with atomic number.

Physical properties

Metallic and ionic radii of actinides

A pellet of ²³⁸PuO₂ to be used in a radioisotope thermoelectric generator for either the Cassini or Galileo mission. The pellet produces 62 watts of heat and glows because of the heat generated by the radioactive decay (primarily α). Photo is taken after insulating the pellet under a graphite blanket for minutes and removing the blanket.

 

Californium

 

Actinides are typical metals. All of them are soft and have a silvery color (but tarnish in air), relatively high density and plasticity. Some of them can be cut with a knife. Their electrical resistivity varies between 15 and 150 µOhm·cm. The hardness of thorium is similar to that of soft steel, so heated pure thorium can be rolled in sheets and pulled into wire. Thorium is nearly half as dense as uranium and plutonium, but is harder than either of them. All actinides are radioactive, paramagnetic, and, with the exception of actinium, have several crystalline phases: plutonium has seven, and uranium, neptunium and californium three. The crystal structures of protactinium, uranium, neptunium and plutonium do not have clear analogs among the lanthanides and are more similar to those of the 3d-transition metals.

All actinides are pyrophoric, especially when finely divided, that is, they spontaneously ignite upon reaction with air. The melting point of actinides does not have a clear dependence on the number of f-electrons. The unusually low melting point of neptunium and plutonium (~640 °C) is explained by hybridization of 5f and 6d orbitals and the formation of directional bonds in these metals.

Chemical properties

Like the lanthanides, all actinides are highly reactive with halogens and chalcogens; however, the actinides react more easily. Actinides, especially those with a small number of 5f-electrons, are prone to hybridization. This is explained by the similarity of the electron energies at the 5f, 7s and 6d shells. Most actinides exhibit a larger variety of valence states, and the most stable are +6 for uranium, +5 for protactinium and neptunium, +4 for thorium and plutonium and +3 for actinium and other actinides.

Chemically, actinium is similar to lanthanum, which is explained by their similar ionic radii and electronic structure. Like lanthanum, actinium almost always has an oxidation state of +3 in compounds, but it is less reactive and has more pronounced basic properties. Among other trivalent actinides Ac³⁺ is least acidic, i.e. has the weakest tendency to hydrolyze in aqueous solutions.

Thorium is rather active chemically. Owing to lack of electrons on 6d and 5f orbitals, the tetravalent thorium compounds are colorless. At pH < 3, the solutions of thorium salts are dominated by the cations [Th(H₂O)₈]⁴⁺. The Th⁴⁺ ion is relatively large, and depending on the coordination number can have a radius between 0.95 and 1.14 Å. As a result, thorium salts have a weak tendency to hydrolyse. The distinctive ability of thorium salts is their high solubility, not only in water, but also in polar organic solvents.

Protactinium exhibits two valence states; the +5 is stable, and the +4 state easily oxidizes to protactinium(V). Thus tetravalent protactinium in solutions is obtained by the action of strong reducing agents in a hydrogen atmosphere. Tetravalent protactinium is chemically similar to uranium(IV) and thorium(IV). Fluorides, phosphates, hypophosphate, iodate and phenylarsonates of protactinium(IV) are insoluble in water and dilute acids. Protactinium forms soluble carbonates. The hydrolytic properties of pentavalent protactinium are close to those of tantalum(V) and niobium(V). The complex chemical behavior of protactinium is a consequence of the start of the filling of the 5f shell in this element.

Uranium has a valence from 3 to 6, the last being most stable. In the hexavalent state, uranium is very similar to the group 6 elements. Many compounds of uranium(IV) and uranium(VI) are non-stoichiometric, i.e. have variable composition. For example, the actual chemical formula of uranium dioxide is UO_2+x, where x varies between −0.4 and 0.32. Uranium(VI) compounds are weak oxidants. Most of them contain the linear "uranyl" group, UO²⁺
₂. Between 4 and 6 ligands can be accommodated in an equatorial plane perpendicular to the uranyl group. The uranyl group acts as a hard acid and forms stronger complexes with oxygen-donor ligands than with nitrogen-donor ligands. NpO²⁺
₂ and PuO²⁺
₂ are also the common form of Np and Pu in the +6 oxidation state. Uranium(IV) compounds exhibit reducing properties, e.g., they are easily oxidized by atmospheric oxygen. Uranium(III) is a very strong reducing agent. Owing to the presence of d-shell, uranium (as well as many other actinides) forms organometallic compounds, such as U^III(C₅H₅)₃ and U^IV(C₅H₅)₄.

Neptunium has valence states from 3 to 7, which can be simultaneously observed in solutions. The most stable state in solution is +5, but the valence +4 is preferred in solid neptunium compounds. Neptunium metal is very reactive. Ions of neptunium are prone to hydrolysis and formation of coordination compounds.

Plutonium also exhibits valence states between 3 and 7 inclusive, and thus is chemically similar to neptunium and uranium. It is highly reactive, and quickly forms an oxide film in air. Plutonium reacts with hydrogen even at temperatures as low as 25–50 °C; it also easily forms halides and intermetallic compounds. Hydrolysis reactions of plutonium ions of different oxidation states are quite diverse. Plutonium(V) can enter polymerization reactions.

The largest chemical diversity among actinides is observed in americium, which can have valence between 2 and 6. Divalent americium is obtained only in dry compounds and non-aqueous solutions (acetonitrile). Oxidation states +3, +5 and +6 are typical for aqueous solutions, but also in the solid state. Tetravalent americium forms stable solid compounds (dioxide, fluoride and hydroxide) as well as complexes in aqueous solutions. It was reported that in alkaline solution americium can be oxidized to the heptavalent state, but these data proved erroneous. The most stable valence of americium is 3 in the aqueous solutions and 3 or 4 in solid compounds.

Valence 3 is dominant in all subsequent elements up to lawrencium (with the exception of nobelium). Curium can be tetravalent in solids (fluoride, dioxide). Berkelium, along with a valence of +3, also shows the valence of +4, more stable than that of curium; the valence 4 is observed in solid fluoride and dioxide. The stability of Bk⁴⁺ in aqueous solution is close to that of Ce⁴⁺.^[92] Only valence 3 was observed for californium, einsteinium and fermium. The divalent state is proven for mendelevium and nobelium, and in nobelium it is more stable than the trivalent state. Lawrencium shows valence 3 both in solutions and solids.

The redox potential ${\displaystyle {\ce {{\mathit {E}}_{\frac {M^{4}+}{AnO2^{2}+}}}}}$ increases from −0.32 V in uranium, through 0.34 V (Np) and 1.04 V (Pu) to 1.34 V in americium revealing the increasing reduction ability of the An⁴⁺ ion from americium to uranium. All actinides form AnH₃ hydrides of black color with salt-like properties. Actinides also produce carbides with the general formula of AnC or AnC₂ (U₂C₃ for uranium) as well as sulfides An₂S₃ and AnS₂.

• 

  Uranyl nitrate (UO₂(NO₃)₂) 
  
  
• 

  Aqueous solutions of uranium III, IV, V, VI salts 
  
  
• 

  Aqueous solutions of neptunium III, IV, V, VI, VII salts 
  
  
• 

  Aqueous solutions of plutonium III, IV, V, VI, VII salts
   
  
• 

  Uranium tetrachloride
   
  
• 

  Uranium hexafluoride
   
  
• 

  U₃O₈ (yellowcake)
  

Compounds

Oxides and hydroxides

Some actinides can exist in several oxide forms such as An₂O₃, AnO₂, An₂O₅ and AnO₃. For all actinides, oxides AnO₃ are amphoteric and An₂O₃, AnO₂ and An₂O₅ are basic, they easily react with water, forming bases:

An₂O₃ + 3 H₂O → 2 An(OH)₃.

These bases are poorly soluble in water and by their activity are close to the hydroxides of rare-earth metals. Np(OH)₃ has not yet been synthesized, Pu(OH)₃ has a blue color while Am(OH)₃ is pink and curium hydroxide Cm(OH)₃ is colorless. Bk(OH)₃ and Cf(OH)₃ are also known, as are tetravalent hydroxides for Np, Pu and Am and pentavalent for Np and Am.

The strongest base is of actinium. All compounds of actinium are colorless, except for black actinium sulfide (Ac₂S₃). Dioxides of tetravalent actinides crystallize in the cubic system, same as in calcium fluoride. 

Thorium reacting with oxygen exclusively forms the dioxide:

${\displaystyle {\ce {Th{}+O2->[{\ce {1000^{\circ }C}}]\overbrace {ThO2} ^{Thorium~dioxide}}}}$

Thorium dioxide is a refractory material with the highest melting point among any known oxide (3390 °C). Adding 0.8–1% ThO₂ to tungsten stabilizes its structure, so the doped filaments have better mechanical stability to vibrations. To dissolve ThO₂ in acids, it is heated to 500–600 °C; heating above 600 °C produces a very resistant to acids and other reagents form of ThO₂. Small addition of fluoride ions catalyses dissolution of thorium dioxide in acids. 

Two protactinium oxides have been obtained: PaO₂ (black) and Pa₂O₅ (white); the former is isomorphic with ThO₂ and the latter is easier to obtain. Both oxides are basic, and Pa(OH)₅ is a weak, poorly soluble base.

Decomposition of certain salts of uranium, for example UO₂(NO₃)·6H₂O in air at 400 °C, yields orange or yellow UO₃. This oxide is amphoteric and forms several hydroxides, the most stable being uranyl hydroxide UO₂(OH)₂. Reaction of uranium(VI) oxide with hydrogen results in uranium dioxide, which is similar in its properties with ThO₂. This oxide is also basic and corresponds to the uranium hydroxide (U(OH)₄).

Plutonium, neptunium and americium form two basic oxides: An₂O₃ and AnO₂. Neptunium trioxide is unstable; thus, only Np₃O₈ could be obtained so far. However, the oxides of plutonium and neptunium with the chemical formula AnO₂ and An₂O₃ are well characterized.

Salts

Einsteinium triiodide glowing in the dark

Actinides easily react with halogens forming salts with the formulas MX₃ and MX₄ (X = halogen). So the first berkelium compound, BkCl₃, was synthesized in 1962 with an amount of 3 nanograms. Like the halogens of rare earth elements, actinide chlorides, bromides, and iodides are water-soluble, and fluorides are insoluble. Uranium easily yields a colorless hexafluoride, which sublimates at a temperature of 56.5 °C; because of its volatility, it is used in the separation of uranium isotopes with gas centrifuge or gaseous diffusion. Actinide hexafluorides have properties close to anhydrides. They are very sensitive to moisture and hydrolyze forming AnO₂F₂. The pentachloride and black hexachloride of uranium were synthesized, but they are both unstable.

Action of acids on actinides yields salts, and if the acids are non-oxidizing then the actinide in the salt is in low-valence state:

U + 2H₂SO₄ → U(SO₄)₂ + 2H₂

2Pu + 6HCl → 2PuCl₃ + 3H₂

However, in these reactions the regenerating hydrogen can react with the metal, forming the corresponding hydride. Uranium reacts with acids and water much more easily than thorium.

Actinide salts can also be obtained by dissolving the corresponding hydroxides in acids. Nitrates, chlorides, sulfates and perchlorates of actinides are water-soluble. When crystallizing from aqueous solutions, these salts forming a hydrates, such as Th(NO₃)₄·6H₂O, Th(SO₄)₂·9H₂O and Pu₂(SO₄)₃·7H₂O. Salts of high-valence actinides easily hydrolyze. So, colorless sulfate, chloride, perchlorate and nitrate of thorium transform into basic salts with formulas Th(OH)₂SO₄ and Th(OH)₃NO₃. The solubility and insolubility of trivalent and tetravalent actinides is like that of lanthanide salts. So phosphates, fluorides, oxalates, iodates and carbonates of actinides are weakly soluble in water; they precipitate as hydrates, such as ThF₄·3H₂O and Th(CrO₄)₂·3H₂O.

Actinides with oxidation state +6, except for the AnO₂²⁺-type cations, form [AnO₄]²⁻, [An₂O₇]²⁻ and other complex anions. For example, uranium, neptunium and plutonium form salts of the Na₂UO₄ (uranate) and (NH₄)₂U₂O₇ (diuranate) types. In comparison with lanthanides, actinides more easily form coordination compounds, and this ability increases with the actinide valence. Trivalent actinides do not form fluoride coordination compounds, whereas tetravalent thorium forms K₂ThF₆, KThF₅, and even K₅ThF₉ complexes. Thorium also forms the corresponding sulfates (for example Na₂SO₄·Th(SO₄)₂·5H₂O), nitrates and thiocyanates. Salts with the general formula An₂Th(NO₃)₆·nH₂O are of coordination nature, with the coordination number of thorium equal to 12. Even easier is to produce complex salts of pentavalent and hexavalent actinides. The most stable coordination compounds of actinides – tetravalent thorium and uranium – are obtained in reactions with diketones, e.g. acetylacetone.

Applications

Interior of a smoke detector containing americium-241.

 

While actinides have some established daily-life applications, such as in smoke detectors (americium) and gas mantles (thorium), they are mostly used in nuclear weapons and use as a fuel in nuclear reactors. The last two areas exploit the property of actinides to release enormous energy in nuclear reactions, which under certain conditions may become self-sustaining chain reaction. 

Self-illumination of a nuclear reactor by Cherenkov radiation.

 

The most important isotope for nuclear power applications is uranium-235. It is used in the thermal reactor, and its concentration in natural uranium does not exceed 0.72%. This isotope strongly absorbs thermal neutrons releasing much energy. One fission act of 1 gram of ²³⁵U converts into about 1 MW·day. Of importance, is that ²³⁵
₉₂U
emits more neutrons than it absorbs; upon reaching the critical mass, ²³⁵
₉₂U
enters into a self-sustaining chain reaction. Typically, uranium nucleus is divided into two fragments with the release of 2–3 neutrons, for example:

²³⁵
₉₂U
+ ¹
₀n
⟶ ¹¹⁵
₄₅Rh
+ ¹¹⁸
₄₇Ag
+ 3¹
₀n

Other promising actinide isotopes for nuclear power are thorium-232 and its product from the thorium fuel cycle, uranium-233. 

Nuclear reactor

The core of most Generation II nuclear reactors contains a set of hollow metal rods, usually made of zirconium alloys, filled with solid nuclear fuel pellets – mostly oxide, carbide, nitride or monosulfide of uranium, plutonium or thorium, or their mixture (the so-called MOX fuel). The most common fuel is oxide of uranium-235.    Fast neutrons are slowed by moderators, which contain water, carbon, deuterium, or beryllium, as thermal neutrons to increase the efficiency of their interaction with uranium-235. The rate of nuclear reaction is controlled by introducing additional rods made of boron or cadmium or a liquid absorbent, usually boric acid. Reactors for plutonium production are called breeder reactor or breeders; they have a different design and use fast neutrons.

Emission of neutrons during the fission of uranium is important not only for maintaining the nuclear chain reaction, but also for the synthesis of the heavier actinides. Uranium-239 converts via β-decay into plutonium-239, which, like uranium-235, is capable of spontaneous fission. The world's first nuclear reactors were built not for energy, but for producing plutonium-239 for nuclear weapons.

About half of the produced thorium is used as the light-emitting material of gas mantles. Thorium is also added into multicomponent alloys of magnesium and zinc. So the Mg-Th alloys are light and strong, but also have high melting point and ductility and thus are widely used in the aviation industry and in the production of missiles. Thorium also has good electron emission properties, with long lifetime and low potential barrier for the emission. The relative content of thorium and uranium isotopes is widely used to estimate the age of various objects, including stars (see radiometric dating).

The major application of plutonium has been in nuclear weapons, where the isotope plutonium-239 was a key component due to its ease of fission and availability. Plutonium-based designs allow reducing the critical mass to about a third of that for uranium-235. The "Fat Man"-type plutonium bombs produced during the Manhattan Project used explosive compression of plutonium to obtain significantly higher densities than normal, combined with a central neutron source to begin the reaction and increase efficiency. Thus only 6.2 kg of plutonium was needed for an explosive yield equivalent to 20 kilotons of TNT. Hypothetically, as little as 4 kg of plutonium—and maybe even less—could be used to make a single atomic bomb using very sophisticated assembly designs.

Plutonium-238 is potentially more efficient isotope for nuclear reactors, since it has smaller critical mass than uranium-235, but it continues to release much thermal energy (0.56 W/g) by decay even when the fission chain reaction is stopped by control rods. Its application is limited by the high price (about US$1000/g). This isotope has been used in thermopiles and water distillation systems of some space satellites and stations. So Galileo and Apollo spacecraft (e.g. Apollo 14) had heaters powered by kilogram quantities of plutonium-238 oxide; this heat is also transformed into electricity with thermopiles. The decay of plutonium-238 produces relatively harmless alpha particles and is not accompanied by gamma-irradiation. Therefore, this isotope (~160 mg) is used as the energy source in heart pacemakers where it lasts about 5 times longer than conventional batteries.

Actinium-227 is used as a neutron source. Its high specific energy (14.5 W/g) and the possibility of obtaining significant quantities of thermally stable compounds are attractive for use in long-lasting thermoelectric generators for remote use. ²²⁸Ac is used as an indicator of radioactivity in chemical research, as it emits high-energy electrons (2.18 MeV) that can be easily detected. ²²⁸Ac-²²⁸Ra mixtures are widely used as an intense gamma-source in industry and medicine.

Development of self-glowing actinide-doped materials with durable crystalline matrices is a new area of actinide utilization as the addition of alpha-emitting radionuclides to some glasses and crystals may confer luminescence.

Toxicity

Schematic illustration of penetration of radiation through sheets of paper, aluminium and lead brick

 

Periodic table with elements colored according to the half-life of their most stable isotope.

  Elements which contain at least one stable isotope.

  Slightly radioactive elements: the most stable isotope is very long-lived, with a half-life of over two million years.

  Significantly radioactive elements: the most stable isotope has half-life between 800 and 34,000 years.

  Radioactive elements: the most stable isotope has half-life between one day and 130 years.

  Highly radioactive elements: the most stable isotope has half-life between several minutes and one day.

  Extremely radioactive elements: the most stable isotope has half-life less than several minutes.

Radioactive substances can harm human health via (i) local skin contamination, (ii) internal exposure due to ingestion of radioactive isotopes, and (iii) external overexposure by β-activity and γ-radiation. Together with radium and transuranium elements, actinium is one of the most dangerous radioactive poisons with high specific α-activity. The most important feature of actinium is its ability to accumulate and remain in the surface layer of skeletons. At the initial stage of poisoning, actinium accumulates in the liver. Another danger of actinium is that it undergoes radioactive decay faster than being excreted. Adsorption from the digestive tract is much smaller (~0.05%) for actinium than radium.

Protactinium in the body tends to accumulate in the kidneys and bones. The maximum safe dose of protactinium in the human body is 0.03 µCi that corresponds to 0.5 micrograms of ²³¹Pa. This isotope, which might be present in the air as aerosol, is 2.5×10⁸ times more toxic than hydrocyanic acid.

Plutonium, when entering the body through air, food or blood (e.g. a wound), mostly settles in the lungs, liver and bones with only about 10% going to other organs, and remains there for decades. The long residence time of plutonium in the body is partly explained by its poor solubility in water. Some isotopes of plutonium emit ionizing α-radiation, which damages the surrounding cells. The median lethal dose (LD₅₀) for 30 days in dogs after intravenous injection of plutonium is 0.32 milligram per kg of body mass, and thus the lethal dose for humans is approximately 22 mg for a person weighing 70 kg; the amount for respiratory exposure should be approximately four times greater. Another estimate assumes that plutonium is 50 times less toxic than radium, and thus permissible content of plutonium in the body should be 5 µg or 0.3 µCi. Such amount is nearly invisible in under microscope. After trials on animals, this maximum permissible dose was reduced to 0.65 µg or 0.04 µCi. Studies on animals also revealed that the most dangerous plutonium exposure route is through inhalation, after which 5–25% of inhaled substances is retained in the body. Depending on the particle size and solubility of the plutonium compounds, plutonium is localized either in the lungs or in the lymphatic system, or is absorbed in the blood and then transported to the liver and bones. Contamination via food is the least likely way. In this case, only about 0.05% of soluble 0.01% insoluble compounds of plutonium absorbs into blood, and the rest is excreted. Exposure of damaged skin to plutonium would retain nearly 100% of it.

Using actinides in nuclear fuel, sealed radioactive sources or advanced materials such as self-glowing crystals has many potential benefits. However, a serious concern is the extremely high radiotoxicity of actinides and their migration in the environment. Use of chemically unstable forms of actinides in MOX and sealed radioactive sources is not appropriate by modern safety standards. There is a challenge to develop stable and durable actinide-bearing materials, which provide safe storage, use and final disposal. A key need is application of actinide solid solutions in durable crystalline host phases.