Actinides
ACTINIDES
CONCEPT
The actinides (sometimes called actinoids) occupy the "bottom line" of the periodic table—a row of elements normally separated from the others, placed at the foot of the chart along with the lanthanides. Both of these families exhibit unusual atomic characteristics, properties that set them apart from the normal sequence on the periodic table. But there is more that distinguishes the actinides, a group of 14 elements along with the transition metal actinium. Only four of them occur in nature, while the other 10 have been produced in laboratories. These 10 are classified, along with the nine elements to the right of actinium on Period 7 of the periodic table, as transuranium (beyond uranium) elements. Few of these elements have important applications in daily life; on the other hand, some of the lower-number transuranium elements do have specialized uses. Likewise several of the naturally occurring actinides are used in areas ranging from medical imaging to powering spacecraft. Then there is uranium, "star" of the actinide series: for centuries it seemed virtually useless; then, in a matter of years, it became the most talked-about element on Earth.
HOW IT WORKS
The Transition Metals
Why are actinides and lanthanides set apart from the periodic table? This can best be explained by reference to the transition metals and their characteristics. Actinides and lanthanides are referred to as inner transition metals, because, although they belong to this larger family, they are usually considered separately—rather like grown children who have married and started families of their own.
The qualities that distinguish the transition metals from the representative or main-group elements on the periodic table are explained in depth within the Transition Metals essay. The reader is encouraged to consult that essay, as well as the one on Families of Elements, which further places the transition metals within the larger context of the periodic table. Here these specifics will be discussed only briefly.
ORBITAL PATTERNS OF THE TRANSITION METALS.
The transition metals are distinguished by their configuration of valence electrons, or the outer-shell electrons involved in chemical bonding. Together with the core electrons, which are at lower energy levels, valence electrons move in areas of probability referred to as orbitals. The pattern of orbitals is determined by the principal energy level of the atom, which indicates a distance that an electron may move away from the nucleus.
Each principal energy level is divided into sublevels corresponding to the number n of the principal energy level. The actinides, which would be on Period 7 if they were included on the periodic table with the other transition metals, have seven principal energy levels. (Note that period number and principal energy level number are the same.) In the seventh principal energy level, there are seven possible sublevels.
The higher the energy level, the larger the number of possible orbital patterns, and the more complex the patterns. Orbital patterns loosely define the overall shape of the electron cloud, but this does not necessarily define the paths along which the electrons move. Rather, it means that if you could take millions of photographs of the electron during a period of a few seconds, the resulting blur of images would describe more or less the shape of a specified orbital.
The four basic types of orbital patterns are discussed in the Transition Metals essay, and will not be presented in any detail here. It is important only to know that, unlike the representative elements, transition metals fill the sublevel corresponding to the d orbitals. In addition, they are the only elements that have valence electrons on two different principal energy levels.
LANTHANIDES AND ACTINIDES.
The lanthanides and actinides are further set apart even from the transition metals, due to the fact that these elements also fill the highly complex f orbitals. Thus these two families are listed by themselves. In most versions of the periodic table, lanthanum (57) is followed by hafnium (72) in the transition metals section of the chart; similarly, actinium (89) is followed by rutherfordium (104). The "missing" metals—lanthanides and actinides, respectively—are shown at the bottom of the chart.
The lanthanides can be defined as those metals that fill the 4f orbital. However, because lanthanum (which does not fill the 4f orbital) exhibits similar properties, it is usually included with the lanthanides. Likewise the actinides can be defined as those metals that fill the 5f orbital; but again, because actinium exhibits similar properties, it is usually included with the actinides.
Isotopes
One of the distinguishing factors in the actinide family is its great number of radioactive isotopes. Two atoms may have the same number of protons, and thus be of the same element, yet differ in their number of neutrons—neutrally charged patterns alongside the protons at the nucleus. Such atoms are called isotopes, atoms of the same element having different masses.
Isotopes are represented symbolically in one of several ways. For instance, there is this format: where S is the chemical symbol of the element, a is the atomic number (the number of protons in its nucleus), and m the mass number—the sum of protons and neutrons. For the isotope known as uranium-238, for instance, this is shown as.
Because the atomic number of any element is established, however, isotopes are usually represented simply with the mass number thus: 238U. They may also be designated with a subscript notation indicating the number of neutrons, so that this information can be obtained at a glance without it being necessary to do the arithmetic. For the uranium isotope shown here, this is written as
Radioactivity
The term radioactivity describes a phenomenon whereby certain materials are subject to a form of decay brought about by the emission of high-energy particles, or radiation.
Types of particles emitted in radiation include:
- Alpha particles, or helium nuclei;
- Beta particles—either an electron or a subatomic particle called a positron;
- Gamma rays or other very high-energy electromagnetic waves.
Isotopes are either stable or unstable, with the unstable variety, known as radioisotopes, being subject to radioactive decay. In this context, "decay" does not mean "rot"; rather, a radioisotope decays by turning into another isotope. By continuing to emit particles, the isotope of one element may even turn into the isotope of another element.
Eventually the radioisotope becomes a stable isotope, one that is not subject to radioactive decay. This is a process that may take seconds, minutes, hours, days, years—and sometimes millions or even billions of years. The rate of decay is gauged by the half-life of a radioisotope sample: in other words, the amount of time it takes for half the nuclei (plural of nucleus) in the sample to become stable.
Actinides decay by a process that begins with what is known as K-capture, in which an electron of a radioactive atom is captured by the nucleus and taken into it. This is followed by the splitting, or fission, of the atom's nucleus. This fission produces enormous amounts of energy, as well as the release of two or more neutrons, which may in turn bring about further K-capture. This is called a chain reaction.
REAL-LIFE APPLICATIONS
The First Three Naturally Occurring Actinides
In the discussion of the actinides that follows, atomic number and chemical symbol will follow the first mention of an element. Atomic mass figures are available on any periodic table, and these will not be mentioned in most cases. The atomic mass figures for actinide elements are very high, as fits their high atomic number, but for most of these, figures are usually for the most stable isotope, which may exist for only a matter of seconds.
Though it gives its name to the group as a whole, actinium (Ac, 89) is not a particularly significant element. Discovered in 1902 by German chemist Friedrich Otto Giesel (1852-1927), it is found in uranium ores. Actinium is 150 times more radioactive than radium, a highly radioactive alkaline earth metal isolated around the same time by French-Polish physicist and chemist Marie Curie (1867-1934) and her husband Pierre (1859-1906).
THORIUM.
More significant than actinium is thorium (Th, 90), first detected in 1815 by the renowned Swedish chemist Jons Berzelius (1779-1848). Berzelius promptly named the element after the Norse god Thor, but eventually concluded that what he had believed to be a new element was actually the compound yttrium phosphate. In 1829, however, he examined another mineral and indeed found the element he believed he had discovered 14 years earlier.
It 1898, Marie Curie and an English chemist named Gerhard Schmidt, working independently, announced that thorium was radioactive. Today it is believed that the enormous amounts of energy released by the radioactive decay of subterranean thorium and uranium plays a significant part in Earth's high internal temperature. The energy stored in the planet's thorium reserves may well be greater than all the energy available from conventional fossil and nuclear fuels combined.
Thorium appears on Earth in an abundance of 15 parts per million (ppm), many times greater than the abundance of uranium. With its high energy levels, thorium has enormous potential as a nuclear fuel. When struck by neutrons, thorium-232 converts to uranium-233, one of the few known fissionable isotopes—that is, isotopes that can be split to start nuclear reactions.
It is perhaps ironic that this element, with its potential for use in some of the most high-tech applications imaginable, is widely applied in a very low-tech fashion. In portable gas lanterns for camping and other situations without electric power, the mantle often contains oxides of thorium and cerium, which, when heated, emit a brilliant white light. Thorium is also used in the manufacture of high-quality glass, and as a catalyst in various industrial processes.
PROTACTINIUM.
Russian chemist Dmitri Ivanovitch Mendeleev (1834-1907), father of the periodic table, used the table's arrangement of elements as a means of predicting the discovery of new substances: wherever he found a "hole" in the table, Mendeleev could say with assurance that a new element would eventually be found to fill it. In 1871, Mendeleev predicted the discovery of "eka-tantalum," an element that filled the space below the transition metal tantalum. (At this point in history, just two years after Mendeleev created the periodic table, the lanthanides and actinides had not been separated from the rest of the elements on the chart.)
Forty years after Mendeleev foretold its existence, two German chemists found what they thought might be Element 91. It had a half-life of only 1.175 minutes, and, for this reason, they named it "brevium." Then in 1918, Austrian physicist Lise Meitner (1878-1968)—who, along with Curie and French physicist Marguerite Perey (1909-1975) was one of several women involved in the discovery of radioactive elements—was working with German chemist Otto Hahn (1879-1968) when the two discovered another isotope of Element 91. This one had a much, much longer half-life: 3.25 · 104 years, or about five times as long as the entire span of human civilization.
Originally named protoactinium, the name of Element 91 was changed to protactinium (Pa), whose longest-lived isotope has an atomic mass of 231. Shiny and malleable, protactinium has a melting point of 2,861.6°F (1,572°C). It is highly toxic, and so rare that no commercial uses have been found for it. Indeed, protactinium could only be produced from the decay products of uranium and radium, and thus it is one of the few elements with an atomic number less than 93 that cannot be said to occur in nature.
Uranium
URANIUM'S EARLY HISTORY.
Chemistry books, in fact, differ as to the number of naturally occurring elements. Some say 88, which is the most correct figure, because protactinium, along with technetium (43) and two others, cannot be said to appear naturally on Earth. Other books say 92, a less accurate figure that nonetheless reflects an indisputable fact: above uranium (U, 92) on the periodic table, there are no elements that generally occur in nature. (However, a few do occur as radioactive by-products of uranium.)
But uranium is much more than just the last truly natural element, though for about a century it apparently had no greater importance. When German chemist Martin Heinrich Klaproth (1743-1817) discovered it in 1789, he named it after another recent discovery: the planet Uranus. During the next 107 years, uranium had a very quiet existence, befitting a rather dull-looking material.
Though it is silvery white when freshly cut, uranium soon develops a thin coating of black uranium oxide, which turns it a flat gray. Yet glassmakers did at least manage to find a use for it—as a coating for decorative glass, to which it imparted a hazy, fluorescent yellowish green hue. Little did they know that they were using one of the most potentially dangerous substances on Earth.
THE DESTRUCTIVE POWER OF URANIUM.
In 1895, German physicist Wilhelm Röntgen (1845-1923) noticed that photographic plates held near a Crookes tube—a device for analyzing electromagnetic radiation—became fogged. He dubbed the rays that had caused this x rays. A year later, in 1896, French physicist Henri Becquerel (1852-1908) left some photographic plates in a drawer with a sample of uranium, and discovered that the uranium likewise caused a fogging of the photographic plates. This meant that uranium was radioactive.
With the development of nuclear fission by Hahn and German chemist Fritz Strassman in 1938, uranium suddenly became all-important because of its ability to undergo nuclear fission, accompanied by the release of huge amounts of energy. During World War II, in what was known as the Manhattan Project, a team of scientists in Los Alamos, New Mexico, developed the atomic bomb. The first of the two atomic bombs dropped on Japan in 1945 contained uranium, while the second contained the transuranium element plutonium.
OTHER USES FOR URANIUM.
Though nuclear weapons have fortunately not been used against human beings since 1945, uranium has remained an important component ofnuclear energy—both in the development of bombs and in the peaceful application of nuclear power. It has other uses as well, due to the fact that it is extremely dense.
Indeed, uranium has a density close to that of gold and platinum, but is much cheaper, because it is more abundant on Earth. In addition, various isotopes of uranium are a by-product of nuclear power, which separates these isotopes from the highly fissionable 235U isotope. Thus, quantities of uranium are available for use in situations where a great deal of mass is required in a small space: in counterweights for aircraft control systems, for instance, or as ballast for missile reentry vehicles.
Because 238U has a very long half-life—4.47 · 109 years, or approximately the age of Earth—it is used to estimate the age of rocks and other geological features. Uranium-238 is the "parent" of a series of "daughter" isotopes that geologists find in uranium ores. Uranium-235 also produces "daughter" isotopes, including isotopes of radium, radon, and other radioactive series. Eventually, uranium isotopes turn into lead, but this can take a very long time: even 235U, which lasts for a much shorter period than 238U, has a half-life of about 700 million years.
The radiation associated with various isotopes of uranium, as well as other radioactive materials, is extremely harmful. It can cause all manner of diseases and birth defects, and is potentially fatal. The tiny amounts of radiation produced by uranium in old pieces of decorative glass is probably not enough to cause any real harm, but the radioactive fallout from Hiroshima resulted in birth defects among the Japanese population during the late 1940s.
Transuranium Elements
Transuranium elements are those elements with atomic numbers higher than that of uranium. None of these occur in nature, except as isotopes that develop in trace amounts in uranium ore. The first such element was neptunium (Np, 93), created in 1940 by American physicist Edwin Mattison McMillan (1907-1991) and American physical chemist Philip Hauge Abelson.
The development of the cyclotron by American physicist Ernest Lawrence (1901-1958) at the University of California at Berkeley in the 1930s made possible the artificial creation of new elements. A cyclotron speeds up protons or ions (charged atoms) and shoots them at atoms of uranium or other elements with the aim of adding positive charges to the nucleus. In the first two decades after the use of the cyclotron to create neptunium, scientists were able to develop eight more elements, all the way up to mendelevium (Md, 101), named in honor of the man who created the periodic table.
Most of these efforts occurred at Berkeley under the leadership of American nuclear chemist Glenn T. Seaborg (1912-1999) and American physicist Albert Ghiorso. The pace of development in transuranium elements slowed after about 1955, however, primarily due to the need for ever more powerful cyclotrons and ion-accelerating machines. In addition to Berkeley, there are two other centers for studying these high-energy elements: the Joint Institute for Nuclear Research in Dubna, Russia, and the Gesellschaft für Schwerionenforschung (GSI) in Darmstadt, Germany.
The Transuranium Actinides
PLUTONIUM.
Just as neptunium had been named for the next planet beyond Uranus, a second transuranium element, discovered by Seaborg and two colleagues in 1940, was named after Pluto. Among the isotopes of plutonium (Pu, 94) is plutonium-239, one of the few fissionable isotopes other than uranium-233 and uranium-235. For that reason, it was applied in the second bomb dropped on Japan.
In addition to its application in nuclear weapons, plutonium is used in nuclear power reactors, and in thermoelectric generators, which convert the heat energy it releases into electricity. Plutonium is also used as a power source in artificial heart pacemakers. Huge amounts of the element are produced each year as a by-product of nuclear power reactors.
BERKELEY DISCOVERIES OF THE 1950S.
Because the lanthanide element above it on the periodic table was named europium after Europe, americium (Am, 95) was named after America. Discovered by Seaborg, Ghiorso, and two others in 1944, it was first produced in a nuclear reaction involving plutonium-239. Americium radiation is used in measuring the thickness of glass during production; in addition, the isotope americium-241 is used as an ionization source in smoke detectors, and in portable devices for taking gamma-ray photographs.
Above curium (Cm, 96) on the periodic table is the lanthanide gadolinium, named after Finnish chemist Johan Gadolin (1760-1852). Therefore, the discoverers of Element 96 alsodecided to name it after a person, Marie Curie. As with some of the other relatively low-number transuranium elements, this one is not entirelyartificial: its most stable isotope, some geologistsbelieve, may have been present in rocks manymillions of years ago, but these isotopes havelong since decayed. Because curium generatesgreat amounts of energy as it decays, it is used forproviding compact sources of power in remotelocations on Earth and in space vehicles.
When Seaborg, Ghiorso, and others created Element 97, berkelium (Bk), they again took the naming of lanthanides as their cue. Just as terbium, directly above it on the periodic table, had been named for the Swedish town of Ytterby, where so many lanthanides were discovered, they named the new element after the American city where so many transuranium elements had been developed. Berkelium has no known applications outside of research. The Berkeley team likewise named californium (Cf, 98) after the state where Berkeley is located. Researchers today are studying the use of californium radiation for treatment of tumors involved in various forms of cancer.
THE REMAINING TRANSURANIUM ACTINIDES.
The remaining transuranium actinides were all named after famous people: einsteinium (Es, 99) for Albert Einstein (1879-1955); fermium (Fm, 100) after Italian-American physicist Enrico Fermi (1901-1954); mendelevium after Mendeleev; nobelium (No, 102) after Swedish inventor and philanthropist Alfred Nobel (1833-1896); and lawrencium (Lr, 103) after Ernest Lawrence.
Both einsteinium and fermium were byproducts of nuclear testing at Bikini Atoll in the south Pacific in 1952. For this reason, their existence was kept a secret for two years. Neither element has a known application. The same is true of mendelevium, produced by Seaborg, Ghiorso, and others with a cyclotron at Berkeley in 1955, as well as the other two transuranium actinides.
Beyond the Transuranium Actinides
As noted earlier, there are nine additional transuranium elements, which properly belong to the transition metals. The first of these is rutherfordium, discovered in 1964 by the Dubna team and named after Ernest Rutherford (1871-1937), the British physicist who discovered the nucleus. The Dubna team named dubnium, discovered in 1967, after their city, just as berkelium had been named after the Berkeley team's city. Both groups developed versions of Element 106 in 1974, and both agreed to name it seaborgium after Seaborg, but this resulted in a controversy that was not settled for some time.
The name of bohrium, created at Dubna in 1976, honors Danish physicist Niels Bohr (1885-1962), who developed much of the model of electron energy levels discussed earlier in this essay. Hassium, produced at the GSI in 1984, is named for the German state of Hess. Two years earlier, the GSI team also created the last named element on the periodic table, meitnerium (109), named after Meitner. Beyond meitnerium are three elements, as yet unnamed, created at the GSI in the mid-1990s.
WHERE TO LEARN MORE
Cooper, Dan. Enrico Fermi and the Revolutions in Modern Physics. New York: Oxford University Press, 1999.
"Exploring the Table of Isotopes" (Web site). <http://ie.lbl.gov/education/isotopes.htm> (May 15, 2001).
Kidd, J. S. and Renee A. Kidd. Quarks and Sparks: The Story of Nuclear Power. New York: Facts on File, 1999.
Knapp, Brian J. Elements. Illustrated by David Woodroffe and David Hardy. Danbury, CT: Grolier Educational, 1996.
"A Periodic Table of the Elements" Los Alamos National Laboratory (Web site). <http://pearl1.lanl.gov/periodic/ (May 22, 2001).
"The Pictorial Periodic Table" (Web site). <http://chemlab.pc.maricopa.edu/periodic/periodic.html> (May 22, 2001).
Sherrow, Victoria. The Making of the Atom Bomb. San Diego, CA: Lucent Books, 2000.
"Some Physics of Uranium." The Uranium Institute (Web site). <http://www.uilondon.org/education_resources/physics_of_uranium/in dex1.htm> (May 27, 2001).
Stwertka, Albert. A Guide to the Elements. New York: Oxford University Press, 1998.
Uranium Information Center (Web site). <http://www.uic.com.au/> (May 27, 2001).
KEY TERMS
ACTINIDES:
Those elements that fill the 5f orbital. Because actinium—which does not fill the 5f orbital—exhibits characteristics similar to those of the actinides, it is usually considered part of the actinidesfamily. Of the other 14 actinides, usually shown at the bottom of the periodic table, only the first three occur in nature.
ATOMIC NUMBER:
The number of protons in the nucleus of an atom. Since this number is different for each element, elements are listed on the periodic table of elements in order of atomic number.
ELECTRON CLOUD:
A term used to describe the pattern formed by orbitals.
HALF-LIFE:
The length of time it takes a substance to diminish to one-half its initialamount. For a sample of radioisotopes, the half-life is the amount of time it takes for half of the nuclei to become stable isotopes. Half-life can be a few seconds; on the other hand, for uranium-238, it is a matter of several billion years.
INNER TRANSITION METALS:
The lanthanides and actinides, which are unique in that they fill the f orbitals. For this reason, they are usually treated separately.
ISOTOPES:
Atoms that have an equal number of protons, and hence are of the same element, but differ in their number of neutrons. This results in a difference ofmass. Isotopes may be either stable or unstable. The unstable type, known as radioisotopes, are radioactive.
LANTHANIDES:
The transition metalsthat fill the 4f orbital.
MASS NUMBER:
The sum of protons and neutrons in the atom's nucleus. The designation 238U, for uranium-238, means that this particular isotope of uranium has a mass number of 238. Since uranium has an atomic number of 92, this means that uranium-238 has 146 neutrons in itsnucleus.
NEUTRON:
A subatomic particle that has no electric charge. Neutrons are found at the nucleus of an atom, alongside protons.
NUCLEUS:
The center of an atom, a region where protons and neutrons are located, and around which electrons spin. The plural of "nucleus" is nuclei.
ORBITAL:
A region of probabilities regarding the position of an electron for anatom in a particular energy state. The higher the principal energy level, the more complex the pattern of orbitals.
PRINCIPAL ENERGY LEVEL:
A value indicating the distance that an electron may move away from the nucleus of anatom. This is designated by a whole-number integer, beginning with 1 and moving upward. The higher the principal energy level, the greater the energy in the atom, and the more complex the pattern of orbitals. Elements in the transition metal family have principal energy levels of 4, 5,6, or 7.
RADIATION:
In a general sense, radiation can refer to anything that travels in astream, whether that stream be composed of subatomic particles or electromagnetic waves. In a more specific sense, the term relates to the radiation from radio active materials, which can be harmful to humanbeings.
RADIOACTIVITY:
A term describing a phenomenon whereby certain isotopes, known as radioisotopes, are subject to a form of decay brought about by the emission of high-energy particles. "Decay" does not mean that the isotope "rots"; rather, it decays to form another isotope until eventually (though this may take a long time) it becomes stable.
RADIOISOTOPE:
An isotope subject to the decay associated with radioactivity. A radioisotope is thus an unstable isotope.
SUBLEVEL:
A region within the principal energy level occupied by electrons in anatom. Whatever the number n of the principal energy level, there are n sublevels. Actinides are distinguished by the fact that their valence electrons are in a sublevel corresponding to the 5f orbital.
TRANSITION METALS:
A group of 40 elements (counting lanthanum and actinium), which are the only elements that fill the d orbital. In addition, the transition metals have their valence electrons on two different principal energy levels. Though the lanthanides and actinides are considered inner transition metals, they are usually considered separately.
TRANSURANIUM ELEMENTS:
Elements with an atomic number higher than that of uranium (92). These have all have been produced artificially. The transuranium elements include 11 actinides, as well as 9 transition metals.
VALENCE ELECTRONS:
Electrons that occupy the highest principal energy level in an atom. These are the electrons involved in chemical bonding.
Actinides
Actinides
The actinides (or actinoids) are the chemical elements with atomic numbers between 90 and 109 inclusively. (An atomic number indicates the number of protons in an atom.) Actinides occur between Groups 3 and 4 in Period 7 of the periodic table. All elements in this family are radioactive (that is, they spontaneously release subatomic particles or energy as their nuclei decay). Five actinides have been found in nature: thorium, protoactinium, uranium, neptunium, and plutonium. The other actinides have been produced artificially in nuclear reactors or particle accelerators (atom-smashers).
History
For many years, the list of chemical elements known to scientists ended with number 92, uranium. Scientists were uncertain as to whether elements heavier than uranium would ever be found. Then, in 1940, a remarkable discovery was made while University of California physicists Edwin McMillan (1907–1991) and Philip Abelson (1913– ) were studying nuclear fission. (Nuclear fission is the splitting of an atomic nucleus, a process that releases large amounts of energy. Atomic bombs and nuclear power plants operate on nuclear fission.) During their research, the duo found evidence for the existence of a new element with atomic number 94, two greater than that of uranium.
This new element was the first transuranium (heaver than uranium) element ever discovered. McMillan and Abelson named it neptunium, after the planet Neptune, just as uranium had been named after the planet Uranus. Later in the same year, McMillan and two other colleagues found a second transuranium element, which they named plutonium, after the planet Pluto.
At that point, the race was on to develop more synthetic transuranium elements, but the research process was not easy. The approach was to fire subatomic particles or small atoms, like those of helium, at a very large nucleus by means of a particle accelerator. If the smaller particle could be made to merge with the larger nucleus, a new atom would be produced. Over time, techniques became more and more sophisticated, and ever-heavier elements were created: americium (number 95) and curium (number 96) in 1944; berkelium (number 97) in 1949; californium (number 98) in 1950; einsteinium (number 99) and fermium (number 100) in 1952; mendelevium (number 101) in 1955; nobelium (number 102) in 1958; and lawrencium (number 103) in 1961.
Studies of the actinide elements are among the most ingenious in all of chemistry. In some cases, no more than one or two atoms of a new element have been produced. Yet scientists have been able to study those few atoms well enough to discover basic properties of the elements. These studies are made even more difficult because most actinide isotopes (atoms of a chemical element that are similar but not exactly alike) decay quickly, with half-lives of only a few days or a few minutes. (A half-life is the amount of time required for half of the atoms of a radioactive substance to disintegrate.)
With the discovery of lawrencium, the actinide family of elements is complete. Scientists have also found elements heavier than lawrencium, but these elements belong to the lanthanide family (or rare earth elements).
Uranium
Uranium is a dull gray metallic element with a melting point of 1,135°C (2,075°F) and a boiling point of 4,134°C (7,473°F). It is relatively abundant in Earth's crust, ranking number 47 among the elements. Although perhaps not as well known, it is actually more abundant than more familiar elements such as tin, silver, mercury, and gold. Natural uranium consists of three isotopes of mass numbers 234 (0.005 percent), 235 (0.711 percent), and 238 (99.283 percent). All three isotopes are radioactive.
Properties and uses. By far the most important property of uranium is its radioactivity. Its most abundant isotope, uranium-238, decays by emitting an alpha particle with a half-life of 4.47 × 109 years. (Recall that the half-life of a radioactive element is the time it takes for one-half of a given sample to decay.) The half-life of uranium-238 is about equal to the age of Earth. That means that about one-half of all the uranium found on Earth at its moment of creation is still here. The other one-half has decayed to other elements.
Knowing the half life of uranium-238 (and many other radioactive isotopes) allows scientists to estimate the age of rocks. The amount of uranium-238 found in any particular rock is compared to the amount of daughter isotopes found with it. A daughter isotope is an isotope formed when some parent isotope, such as uranium-238, decays. The more daughter isotope present in a sample, the older the rock; the less daughter isotope, the younger the rock.
The second most abundant isotope of uranium, uranium-235, has the rare property of being fissionable, meaning that its atomic nuclei will break apart when bombarded by neutrons. The fission of a uranium-235 nucleus releases very large amounts of energy, additional neutrons, and two large fission products. The fission products are the atomic nuclei formed when a fissionable nucleus such as uranium-235 breaks apart.
The fission of uranium-235 nuclei has become extremely important in the manufacture of nuclear weapons and in the operation of nuclear power plants. In fact, these applications account for the primary applications of uranium in everyday life.
Thorium
Thorium is a soft metal with a bright silvery luster when freshly cut. It has a melting point of about 1,700°C (3,100°F) and a boiling point of about 4,500°C (8,100°F). It is relatively soft, with a hardness about equal to that of lead. It is even more abundant than uranium, ranking number 39 in abundance among the elements in Earth's crust.
No more than a few hundred tons of thorium are produced annually. About one-half of this production goes to the manufacture of gas mantles, insulated chambers in which fuel is burned. The rest goes for use as nuclear fuel, in sunlamps (electric lamps that emit radiation; often used for tanning), in photoelectric cells (vacuum tubes in which electric current flows when light strikes the photosensitive—or light sensitive—cathode), and in the production of other alloys (a mixture of two or more metals or a metal and a nonmetal).
Uses of other actinides
At one time, the actinides other than uranium were no more than scientific curiosities. They were fascinating topics of research for scientists but of little practical interest. That situation has now changed, and all of the actinides that can be prepared in large enough quantities have found some use or another. Plutonium, for example, is used in the manufacture of nuclear weapons and as the power source in nuclear power plants. On a smaller scale, it is also used as a power source in smaller devices such as the heart pacemaker. Californium is used in smoke detectors, curium is a power source in space vehicles, and americium is utilized in the treatment of cancer.
[See also Element, chemical; Isotope; Nuclear fission; Periodic table; Radioactivity ]
Actinides
Actinides
Physical and chemical properties
Actinides, or actinoids, is a generic term that refers to a series of 15 radioactive chemical elements with atomic numbers 89 (actinium) through 103 (lawrencium). Denoted by the generic symbol An, these elements are all radioactive heavy metals, positioned in the seventh period and elaborated upon at the bottom of the periodic table. In order of increasing atomic number, the actinide series members are: actinium, thorium, protactinium, uranium, neptunium, plutonium, americium, curium, berkelium, californium, einsteinium, fermium, mendelevium, nobelium, and lawrencium. (Note: Lawrencium is sometimes not grouped with the actinides but with the transition metals.) The actinide elements share the following properties: radioactive; easily tarnished in air; highly electropositive; very dense metals with unique structures; ability to combine directly with most non-metals; and ability to react with dilute acids or boiling water with the subsequent release of hydrogen gas.
Occurrence
Only actinium (atomic symbol, Ac), thorium (Th), protactinium (Pa), and uranium (U) are extracted from deposits in nature. In Canada, the United States, South Africa, and Namibia, thorium and protactinium are available in large quantities. All other actinides are synthetic, or artificially made.
To understand the physical and chemical properties of actinides, a basic foundation of atomic structure, radioactivity, and the periodic table is required. The atomic structure can be pictured like a solar system. In the middle of the atom is the nucleus, composed of neutrons (no charge) and protons (positively charged). Around the nucleus, electrons (negatively charged) are rotating on their own axis, as well as circulating in definite energy levels. Each energy level (or shell) is designated by a principal quantum number (n) as K, L, M, N, O, etc. or 1, 2, 3, 4, 5, etc., respectively. Each shell has sub-shells, or orbitals. The first energy level consists of one orbital (s); the second level consists of two orbitals (s and p); the third level consists of three orbitals (s, p, and d), and from the fourth level on up there are four orbitals (s, p, d, and f). The orbitals closer to the nucleus are lower in energy level than the orbitals further away from the nucleus.
The electrons are distributed according to Pauli’s exclusion principle. In any atom, the number of protons is equal to the number of electrons, thus bringing neutrality in charge. These stable and abundant atoms exist in nature only. In the unstable and less abundant atoms, the number of neutrons is more than the number of electrons (one element with the same atomic number but with a different atomic mass). These unstable atoms are known as isotopes, some of which are radioactive.
Radioactive isotopes become nonradioactive by the decaying process. The decaying process may involve an emission of: (1) electrons or negative beta particles; (2) helium nuclei or alpha particles; (3) gamma rays or very high frequency electromagnetic waves; or (4) positrons or positively charged electrons or positive beta particles.
The decaying process may also be due to K-capture (an orbital electron of a radioactive atom that may be captured by the nucleus and taken into it). Each of the above mentioned decay processes results in an isotope of a different element (an element with a different atomic number). The emission of alpha particles also results in elements with different atomic weights.
The most important decay process in actinides is K-capture, followed by the splitting, or fission, of the nucleus. This fission results in enormous amounts of energy and two or more extra neutrons. These newly formed neutrons can further start K-capture, with the subsequent reactions going on like a chain reaction. Atomic reactors and atomic bombs depend on the chain reactions.
A scheme of the classification of all known (both discovered and artificially-made) elements is represented in the modern periodic table. The periodic table is divided into vertical columns and horizontal rows representative of the periods with increasing atomic numbers. Each box contains one element and is represented by its symbol, a single or double letter, with the atomic number as a superscript, and the atomic weight as a subscript. Note that in the sixth and seventh periods, there are breaks between atomic numbers 57 and 72 (lanthanides) and 89 and 104 (actinides). Fourteen elements are present between the atomic numbers 89 and 104; these are elaborated upon at the bottom of the periodic table. These 14 elements, plus actinium, are known as actinides.
General preparation
All actinide metals are prepared on a scale by the reduction of AnF3 or AnF4 with vapors of lithium (Li), magnesium (Mg), calcium (Ca), or barium (Ba) at 2, 102–2, 552°F (1, 150–1, 400°C); sometimes chlorides or oxides are used. The Van Arkel-de Boer process is a special preparation method used for thorium and protactinium.
Physical and chemical properties
A common feature of actinides is the possession of multiple oxidation states. The term oxidation state refers to the number of electron(s) that are involved or that can possibly become involved in the formation of chemical bond(s) in that compound, when one element combines with another element during a chemical reaction. The oxidation state is designated by a plus sign (when an electron is donated or electro-positive) or by a minus sign (when the electron is accepted or electronegative). An element can have more than one oxidation state. An electron configuration can provide the information about the oxidation state of that element. The most predominant oxidation state among actinides is +3, which is similar to lanthanides. The crystal structure (geometry), solubility property, and the formation of chemical compounds are based on the oxidation state of the given element.
Actinides ions in an aqueous solution are colorful, containing colors such as red purple (U3+), purple (Np3+), pink (Am3+), green (U4+), yellow green (Np4+), and pink red (Am4+). Actinides ions U, Np, Pu, and Am undergo hydrolysis, disproportionation, or formation of polymeric ions in aqueous solutions with a low pH.
All actinides are characterized by partially filled 5f, 6d, and 7s orbitals. Actinides form complexes easily with certain ligands as well as with halide, sulfate, and other ions. Organometallic compounds (compounds with a sign bond between the metal and carbon atom of organic moiety) of uranium and thorium have been prepared and are useful in organic synthesis. Several alloys of protactinium with uranium have been prepared.
Uses of actinides
Even though hazards are associated with radioactivity of actinides, many beneficial applications exist as well. Radioactive nuclides are used in cancer therapy, analytical chemistry, and in basic research in the study of chemical structures and mechanisms. The explosive power of uranium and plutonium are exploited in making atom bombs. In fact, the uranium-enriched atom bomb that exploded over Japan was the first uranium bomb released. Nuclear reactions of uranium-235 and plutonium-239 are currently utilized in atomic energy power plants to generate electric power. Thorium is economically useful for the reason that fissionable uranium-233 can be produced from thorium-232. Plutonium-238 is used in implants in the human body to power the heart pacemaker, which does not need to be replaced for at least 10 years. Curium-244 and plutonium-238 emit heat at 2.9 watts and 0.57 watts per gram, respectively. Therefore, curium and plutonium are used as power sources on the Moon and on spacecraft to provide electrical energy for transmitting messages to Earth.
Sutharchanadevi Murugen
Actinides
Actinides
Actinides or actinoids is a generic term that refers to a series of 15 chemical elements. Denoted by the generic symbol An, these elements are all radioactive heavy metals, positioned in the seventh period and elaborated upon at the bottom of the periodic table .
Occurrence
Only actinium (atomic symbol, Ac), thorium (Th), protactinium (Pa), and uranium (U) are extracted from deposits in nature. In Canada, the United States, South Africa, and Namibia, thorium and protactinium are available in large quantities. All other actinides are synthetic or man-made.
To understand the physical and chemical properties of actinides, a basic foundation of atomic structure, radioactivity, and the periodic table is required. The atomic structure can be pictured like a solar system. In the middle of the atom is the nucleus, composed of neutrons (no charge) and protons (positively charged). Around the nucleus, electrons (negatively charged) are rotating on their own axis, as well as circulating in definite energy levels. Each energy level (or shell) is designated by a principal quantum number (n) as K, L, M, N, O, etc. or 1, 2, 3, 4, 5, etc. respectively. Each shell has sub-shells, or orbitals. The first energy level consists of one orbital (s); the second level consists of two orbitals (s and p); the third level consists of three orbitals (s, p, and d), and from the fourth level on up there are four orbitals (s, p, d, and f). The orbitals closer to the nucleus are lower in energy level than the orbitals further away from the nucleus.
The electrons are distributed according to Pauli's exclusion principle. In any atom, the number of protons is equal to the number of electrons, thus bringing the neutrality in charge. These stable and abundant atoms exist in nature only. In the unstable and less abundant atoms, the number of neutrons is more than the number of electrons (one element with the same atomic number but with a different atomic mass ). These unstable atoms are known as isotopes, some of which are radioactive.
Radioactive isotopes become nonradioactive by the decaying process. The decaying process may involve an emission of: (1) electrons or negative beta particles; (2) helium nuclei or alpha particles; (3) gamma rays or very high frequency electromagnetic waves; (4) positrons or positively charged electrons or positive beta particles.
The decaying process may also be due to K-capture (an orbital electron of a radioactive atom that may be captured by the nucleus and taken into it). Each of the above mentioned decay processes results in a isotope of a different element (an element with a different atomic number). The emission of alpha particles also results in elements with different atomic weights.
The most important decay process in actinides is K-capture, followed by the splitting, or fission, of the nucleus. This fission results in enormous amounts of energy and two or more extra neutrons. These newly formed neutrons can further start K-capture, with the subsequent reactions going on like a chain reaction. Atomic reactors and atomic bombs depend on the chain reactions.
A scheme of the classification of all known (both discovered and man-made) elements is represented in the modern periodic table. The periodical table is divided into vertical columns and horizontal rows representative of the periods with increasing atomic numbers. Each box contains one element and is represented by its symbol, a single or double letter, with the atomic number as a superscript, and the atomic weight as a subscript. Note that in the sixth and seventh periods, there are breaks between atomic numbers 57 and 72 (lanthanides ) and 89 and 104 (actinides). Fourteen elements are present between the atomic numbers 89 and 104, and are elaborated upon at the bottom of the periodic table. These 14 elements, plus actinium, are known as actinides.
General preparation
All actinide metals are prepared on a scale by the reduction of AnF3 or AnF4 with vapors of lithium (Li), magnesium (Mg), calcium (Ca), or barium (Ba) at 2,102–2,552°F (1,150–1,400°C); sometimes chlorides or oxides are used. The Van Arkel-de Boer process is a special preparation method used for thorium and protactinium.
Physical and chemical properties
A common feature of actinides is the possession of multiple oxidation states. The term oxidation state refers to the number of electron(s) that are involved or that can possibly become involved in the formation of chemical bond(s) in that compound, when one element combines with another element during a chemical reaction. The oxidation state is designated by a plus sign (when an electron is donated or electro-positive) or by a minus sign (when the electron is accepted or electronegative). An element can have more than one oxidation state. An electron configuration can provide the information about the oxidation state of that element. The most predominant oxidation state among actinides is +3, which is similar to lanthanides. The crystal structure (geometry ), solubility property, and the formation of chemical compounds are based on the oxidation state of the given element.
Actinides ions in an aqueous solution are colorful, containing colors such as red purple (U3+), purple (Np3+), pink (Am3+), green (U4+), yellow green (Np4+), and pink red (Am4+). Actinides ions U, Np, Pu, and Am undergo hydrolysis , disproportionation, or formation of polymeric ions in aqueous solutions with a low pH .
All actinides are characterized by partially filled 5f, 6d, and 7s orbitals. Actinides form complexes easily with certain ligands as well as with halide, sulfate, and other ions. Organometallic compounds (compounds with a sign bond between the meta and carbon atom of organic moiety) of uranium and thorium have been prepared and are useful in organic synthesis. Several alloys of protactinium with uranium have been prepared.
Uses of actinides
Even though hazards are associated with radioactivity of actinides, many beneficial applications exist as well. Radioactive nuclides are used in cancer therapy, analytical chemistry , and in basic research in the study of chemical structures and mechanisms. The explosive power of uranium and plutonium are well exploited in making atom bombs. In fact, the uranium enriched atom bomb that exploded over Japan was the first uranium bomb released. Nuclear reactions of uranium-235 and plutonium-239 are currently utilized in atomic energy powerplants to generate electric power. Thorium is economically useful for the reason that fissionable uranium-233 can be produced from thorium-232. Plutonium-238 is used in implants in the human body to power the heart pacemaker , which is does not need to be replaced for at least 10 years. Curium-244 and plutonium-238 emit heat at 2.9 watts and 0.57 watts per gram, respectively. Therefore, curium and plutonium are used as power sources on the Moon to provide electrical energy for transmitting messages to Earth .
Sutharchanadevi Murugen
Actinides
Actinides
The actinide elements (atomic numbers 89 through 103) involve the filling of 5f orbitals. All actinides are radioactive, but only uranium and the lighter actinides have half-lives long enough to be present in Earth's environment. The heavier actinides are produced by nuclear reactions and some have very short half-lives. The actinides also undergo a radius contraction as do the lanthanides with an increasing atomic number . They are characterized by variable oxidation numbers, but the importance of the +3 state and the similarities to lanthanides increase for heavier elements. American chemist Glenn Seaborg is credited with revising the Periodic Table so actinides were placed under lanthanides.
see also Actinium; Americium; Berkelium; Einsteinium; Fermium; Lawrencium; Mendelevium; Neptunium; Nobelium; Plutonium; Protactinium; Rutherfordium; Seaborg, Glenn Theodore.
Herbert B. Silber
Bibliography
Cotton, Simon (1991). Lanthanides and Actinides. New York: Oxford University Press.
actinide series
Their atomic numbers range from 89 to 103. Each element is analogous to the corresponding lanthanide series (rare-earth metals). Those having atomic numbers greater than 92 (uranium) are the transuranic elements.