Actinide

Actinide

Actinides in the periodic table

The actinide or actinoid (IUPAC nomenclature) series encompasses the 15 metallic chemical elements with atomic numbers from 89 to 103, actinium through lawrencium.[2][3][4][5]

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, corresponding to the filling of the 5f electron shell; lawrencium, a d-block element, is also generally considered an actinide. 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 through neptunium are much more similar to transition metals in their chemistry.

Of the actinides, primordial thorium and uranium occur naturally in substantial quantities and small amounts of persisting natural plutonium have also been identified. 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.[2][6] 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.[7]

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.

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,[2] 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).

Actinides
Actin­ium
89
Ac
Thor­ium
90
Th
Protac­tinium
91
Pa
Ura­nium
92
U
Neptu­nium
93
Np
Pluto­nium
94
Pu
Ameri­cium
95
Am
Curium
96
Cm
Berkel­ium
97
Bk
Califor­nium
98
Cf
Einstei­nium
99
Es
Fer­mium
100
Fm
Mende­levium
101
Md
Nobel­ium
102
No
Lawren­cium
103
Lr
Primordial From decay Synthetic Border shows natural occurrence of the element

Discovery, isolation and synthesis

Synthesis of transuranium elements[8][9]
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 α-particles
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 most abundant, or easy to synthesize actinides are uranium and thorium, followed by plutonium, americium, actinium, protactinium and neptunium.[10]

The existence of transuranium elements was suggested by [13][14]

At present, there are two major methods of producing isotopes of transplutonium elements: irradiation of the lighter elements with either neutrons or 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.[15]

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 inobservation 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.[16]

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,[6] 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.[17] Only 60 years later, the French scientist Eugène-Melchior Péligot identified it with uranium oxide. He also isolated the first sample of uranium metal by heating uranium tetrachloride with potassium.[18] 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.[19][20]

Thorium oxide was discovered by Friedrich Wöhler in the mineral, which was found in Norway (1827).[21] 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.[22][23] The same isolation method was later used by Péligot for uranium.[6]

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[24] and (in 1900) as similar to thorium.[25] The discovery of actinium by Debierne was however questioned in 1971[26] and 2000,[27] arguing that Debierne's publications in 1904 contradicted his earlier work of 1899–1900. 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.[28][29] 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.[30][31]

Protactinium was possibly isolated in 1900 by William Crookes.[32] 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);[33][34] 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 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. produced 130 grams of protactinium from 60 tonnes of waste left after extraction of uranium from its ore.[35]

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.[36] They produced the 239Np isotope (half-life = 2.4 days) by bombarding uranium with slow neutrons.[35] It was the first transuranium element produced synthetically.[37]

seaborgium in his honor while he was still living. They also synthesized more than 100 atomic 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:

\mathrm{ {}^{238}_{92}U + {}^{1}_{0}n\ \xrightarrow\ \ {}^{239}_{92}U\ \xrightarrow[23.5\ min]{\beta^-}\ {}^{239}_{93}Np\ \xrightarrow[2.3\ days]{\beta^-}\ {}^{239}_{94}Pu\ \xrightarrow[2.4\cdot 10^4\ years]{\alpha} }

In this way, Enrico Fermi with collaborators, using the first nuclear reactor Chicago Pile-1, obtained significant amounts of plutonium-239, which were then used in nuclear weapons.[38]

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

~\mathrm{AnO^{2+}_{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.[85]

Compounds

Oxides and hydroxides

An – actinide
**Depending on the isotopes

Some actinides can exists 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:[85]

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. The strongest base is of actinium. All compounds of actinium are colorless, except for black actinium sulfide (Ac2S3).[85] Dioxides of tetravalent actinides crystallize in the cubic system, same as in calcium fluoride.

Thorium reacting with oxygen exclusively forms dioxide:

~\mathrm{Th+O_2 \ \xrightarrow{1000^\circ C} \ ThO_2}

Thorium dioxide is a refractory material with the highest melting point among any known oxide (3390 °C).[94] 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 were 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.[85]

Decomposition of certain salts of uranium, for example UO2(NO3)·6H2O in air at 400 °C, yields orange or yellow UO3.[94] This oxide is amphoteric and forms several hydroxides, the most stable being 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).[85]

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.[85]

Salts

*An – actinide
**Depending on the isotopes
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.[99] The pentachloride and black hexachloride of uranium were synthesized, but they are both unstable.[85]

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 + 2 H2SO4 → U (SO4)2 + 2 H2
2 Pu + 6 HCl → 2 PuCl3 + 3 H2

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.[85]

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.[85]

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.[85]

Applications

Interior of a smoke detector containing americium-241.

While actinides have some established daily-life applications, such as in smoke detectors (americium)[100][101] and gas mantles (thorium),[75] they are mostly used in nuclear weapons and use as a fuel in nuclear reactors.[75] 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 235U emits more neutrons than it absorbs;[102] upon reaching the critical mass, 235U enters into a self-sustaining chain reaction.[69] Typically, uranium nucleus is divided into two fragments with the release of 2–3 neutrons, for example:

~\mathrm{{}_{~}^{235}U+ \ ^1_0n \longrightarrow {}_{~}^{115}Rh+{}_{~}^{118}Ag + 3^1_0n}
Other promising actinide isotopes for nuclear power are thorium-232 and its product from the thorium fuel cycle, uranium-233.
Nuclear reactor[69][103][104]
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.
Nuclear reactor scheme

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.[75] 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.[102] The relative content of thorium and uranium isotopes is widely used to estimate the age of various objects, including stars (see radiometric dating).[105]

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.[106] 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.[107] (See also Nuclear weapon design.) 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.[108]

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)[101][109] by decay even when the fission chain reaction is stopped by control rods. Its application is limited by the high price (about 1000 USD/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[110]) 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.[101]

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.[29]

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.[111]

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 four million years.
  Moderately radioactive elements: the most stable isotope has half-life between 800 and 34,000 years.
  Very radioactive elements: the most stable isotope has half-life between one day and 103 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. Very little is known about these elements due to their extreme instability and radioactivity.

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.[29]

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.[61]

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.[87]

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.[112] 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.[111]

See also

References and notes

  1. ^ The Manhattan Project. An Interactive History. US Department of Energy
  2. ^ a b c
  3. ^ Actinide element, Encyclopædia Britannica on-line
  4. ^ Although "actinoid" (rather than "actinide") means "actinium-like" and therefore should exclude actinium, that element is usually included in the series.
  5. ^
  6. ^ a b c Greenwood, p. 1250
  7. ^ a b
  8. ^ a b c Greenwood, p. 1252
  9. ^ Nobelium and lawrencium were almost simultaneously discovered by Soviet and American scientists
  10. ^ Myasoedov, p. 7
  11. ^
  12. ^
  13. ^
  14. ^
  15. ^ Myasoedov, p. 9
  16. ^ Myasoedov, p. 14
  17. ^
  18. ^
  19. ^
  20. ^ K. Zimmerman, Ann., 213, 290 (1882); 216, 1 (1883); Ber. 15 (1882) 849
  21. ^ Golub, p. 214
  22. ^ (modern citation: Annalen der Physik, vol. 92, no. 7, pages 385–415)
  23. ^
  24. ^
  25. ^
  26. ^
  27. ^
  28. ^ Golub, p. 213
  29. ^ a b c d e f g h i j
  30. ^
  31. ^
  32. ^
  33. ^ a b
  34. ^
  35. ^ a b Greenwood, p. 1251
  36. ^
  37. ^ a b c d e f
  38. ^
  39. ^
  40. ^ Myasoedov, p. 8
  41. ^
  42. ^
  43. ^ Wallace W. Schulz (1976) The Chemistry of Americium, U.S. Department of Commerce, p. 1
  44. ^
  45. ^
  46. ^
  47. ^
  48. ^
  49. ^ S. G. Thompson, B. B. Cunningham: "First Macroscopic Observations of the Chemical Properties of Berkelium and Californium", supplement to Paper P/825 presented at the Second Intl. Conf., Peaceful Uses Atomic Energy, Geneva, 1958
  50. ^ Darleane C. Hoffman, Albert Ghiorso, Glenn Theodore Seaborg (2000) The transuranium people: the inside story, Imperial College Press, ISBN 1-86094-087-0, pp. 141–142
  51. ^ a b
  52. ^
  53. ^
  54. ^
  55. ^
  56. ^ a b c d e f g
  57. ^ a b c d e f g h i
  58. ^ Myasoedov, pp. 19–21
  59. ^ Specific activity is calculated by given in the table half-lives and the probability of spontaneous fission
  60. ^ a b Greenwood, p. 1254
  61. ^ a b c d e f g
  62. ^
  63. ^ a b Myasoedov, p. 18
  64. ^ a b c Myasoedov, p. 22
  65. ^ Myasoedov, p. 25
  66. ^
  67. ^
  68. ^
  69. ^ a b c d e f g h i
  70. ^
  71. ^ a b c
  72. ^ Thorium, USGS Mineral Commodities
  73. ^ a b c d e f g Golub, pp. 215–217
  74. ^ Greenwood, pp. 1255, 1261
  75. ^ a b c d e Greenwood, p. 1255
  76. ^
  77. ^
  78. ^ Golub, pp. 218–219
  79. ^ a b c Greenwood, p. 1263
  80. ^ a b
  81. ^
  82. ^ Greenwood, p. 1265
  83. ^ a b Greenwood, p. 1264
  84. ^ Myasoedov, pp. 30–31
  85. ^ a b c d e f g h i j k Golub, pp. 222–227
  86. ^ Greenwood, p. 1278
  87. ^ a b
  88. ^
  89. ^ a b Myasoedov, pp. 25–29
  90. ^ Myasoedov, p. 88
  91. ^ a b
  92. ^ According to other sources, cubic sesquioxide of curium is olive-green. See
  93. ^ The atmosphere during the synthesis affects the lattice parameters, which might be due to non-stoichiometry as a result of oxidation or reduction of the trivalent californium. Main form is the cubic oxide of californium(III).
  94. ^ a b c Greenwood, p. 1268
  95. ^
  96. ^ a b Greenwood, p. 1270
  97. ^ Myasoedov, pp. 96–99
  98. ^
  99. ^ Greenwood, p.1269
  100. ^ Smoke Detectors and Americium, Nuclear Issues Briefing Paper 35, May 2002
  101. ^ a b c Greenwood, p. 1262
  102. ^ a b Golub, pp. 220–221
  103. ^
  104. ^ Greenwood, pp. 1256–1261
  105. ^
  106. ^
  107. ^
  108. ^
  109. ^ John Holdren and Matthew Bunn Nuclear Weapons Design & Materials. Project on Managing the Atom (MTA) for NTI. 25 November 2002
  110. ^ Apollo 14 Press Kit – 01/11/71, NASA, pp. 38–39
  111. ^ a b
  112. ^

Bibliography

External links

  • Lawrence Berkeley Laboratory image of historic periodic table by Seaborg showing actinide series for the first time
  • Uncovering the Secrets of the ActinidesLawrence Livermore National Laboratory,
  • Actinide Research QuarterlyLos Alamos National Laboratory,