Periodic Table
Periodic Table
The arrangement of the chemical elements into periods (horizontal rows) and groups (vertical columns) is called the periodic table. The periodic table is an essential part of the language of chemistry. It has much in common with a thesaurus, providing a guide to similarities and differences among the elements. From the way elements are organized in the periodic table, chemists can predict their behavior and write chemical formulas of compounds using just a few general guidelines. Using such rules is not the same as understanding why elements in certain areas the periodic table behave as they do, but the trends that arise from the arrangement of elements in the periodic table allows a chemist to remember useful facts about the types of compounds formed from specific elements and their chemical reactions.
The elements in the table are represented by symbols (one, two, or three letters) in individual squares. Above each chemical symbol appears the atomic number of the element. These whole numbers are the number of protons present in the nucleus of that element. Below the element symbol appears the atomic mass, which is the average mass of all the isotopes of that element. The elements are arranged in order of increasing atomic numbers. Elements of the same group are found to have similar chemical properties.
The ultimate effectiveness of the periodic table is that it arranges over one hundred individual elements so information about a given element is known merely by where it is found in the periodic table. The discovery that elements could be arranged in a periodic table was made by Russian chemist Dmitri Ivanovitch Mendeleev (1834-1907). Since its discovery in 1869, the periodic table has guided chemical research including the discovery of new elements. This ability to lead scientific inquiry over a 137-year span has contributed to the periodic table being considered one of the greatest scientific constructs. The magnitude of the scientific time span over which the periodic table has guided research is more strikingly illustrated when it is considered that it has been used from a time prior to the discovery of the light bulb to a time after the start of the construction of the International Space Station, the operations of the NASA space shuttle fleet, and the early beginnings of the Chinese space program.
Construction of the table
The discovery of the individual elements was a necessary prerequisite for the construction of the periodic table. The first pure elements have been known since the time of the ancient Greeks who used the metallic elements gold (Au), silver (Ag), tin (Sn), copper (Cu), lead (Pb), and mercury (Hg). The first individual credited with the discovery of an element was German alchemist Hennig Brand (c. 1630-1710), a scientist who discovered the element phosphorous in 1649.
Mendeleev was born in Siberia and studied chemistry at St. Petersburg Institute in Russia. He went on to become a science teacher and later a lecturer and researcher at the University of St. Petersburg. It was through his experience as a teacher that Mendeleev realized a classification system of the known elements was needed. Earlier attempts were made to order the known elements, but they suffered from either being too simplistic or led to inconsistencies that limited their usefulness.
In the second half of the nineteenth century, data from laboratories in France, England, Germany, and Italy were assembled into a pamphlet by Italian chemist Stanislao Cannizzaro (1826-1910), a teacher in what is now northern Italy. In this pamphlet, Cannizzaro demonstrated a way to determine a consistent set of atomic weights, one weight for each of the elements then known. Cannizzaro distributed his pamphlet and explained his ideas at an international meeting held in Karlsruhe, Germany, in 1860, which was organized to discuss new ideas about the theory of atoms. When Mendeleyev returned from the meeting to St. Petersburg, Russia, he pondered Cannizzaro’s list of atomic weights along with an immense amount of information he had gathered about the properties of elements. Mendeleyev found that when he arranged the elements in order of increasing atomic weight, similar properties were repeated at regular intervals; they displayed periodicity. He used the periodic repetition of chemical and physical properties to construct a chart much like the periodic table currently in use.
With the purpose of assembling the 63 known elements into an ordered system, Mendeleev wrote the elements’ names on individual cards. The cards also contained the atomic mass and specific gravity as well as other known chemical data for that element. By arranging the cards by increasing atomic mass in rows, and then in columns so that elements having similar chemical properties would lie under each other, the first periodic table was formed.
Mendeleev’s predictions
Mendeleev came to believe in his periodic table to such a degree that he changed the atomic mass of known elements so that they fit where they belonged in his table. He did this with no experimental evidence, only his belief in his table. In one such case he changed the atomic mass of beryllium (Be) from 14, which placed it in Group 15 above nitrogen (N), to 9. This placed it in Group 2 above magnesium (Mg) with which it was more closely related chemically. Even more daring was the fact that Mendeleev predicted the properties of undiscovered elements. Based on gaps in the periodic table Mendeleev deduced that in these gaps belonged elements yet to be discovered. Based on other elements in the same group he predicted the existence of eka-aluminum, eka-boron, and eka-silicon, later to be named gallium (Ga), scandium (Sc), and germanium (Ge), respectively. Mendeleev predicted the atomic mass of each element along with compounds they each should form. Within fifteen years of Mendeleev’s predictions, these elements were discovered, and their properties were found to closely match his predictions. These fulfilled predictions went a long way toward convincing any remaining doubters of the infallibility of the periodic table.
Another change Mendeleev made based on chemical analogy and intuition was placing iodine (I) after tellurium (Te), even though the atomic mass of iodine was less than tellurium. This anomaly, along with the difficulty of where to place the inner transition metals, was a problem that would soon be definitively solved. At the time of the periodic table’s construction, little was known of atomic structure. With further scientific discoveries such as the existence of protons and the existence of electronic shells, these mysteries were explained and placed into their current places in the periodic table.
Early in the twentieth century, work initiated by English physicist Joseph John Thomson (1856-1940) led to the discovery of the electron and, later, the proton. In 1932, English physicist James Chadwick (1891-1974) proved the existence of the neutron in the atomic nucleus. The discovery of these elementary particles and the experimental determination of their actual weights led scientists to conclude that different atoms have different weights because they contain different numbers of protons and neutrons. However, it was not yet clear how many subatomic particles were present in any but the simplest atoms, such as hydrogen, helium, and lithium.
Refinement in the measurements of atomic mass, the ordering of the elements based on atomic number rather than atomic mass by British scientist Henry G.J. Moseley 1887-1915) in 1913, and the discovery of new elements have led to the continuing evolution of the periodic table. Moseley determined the frequency and wavelength of x rays emitted by a large number of elements. By this time, the number of protons in the nucleus of some of the lighter elements had been determined. Moseley found the wavelength of the most energetic x ray of an element decreased systematically as the number of protons in the nucleus increased. Moseley then hypothesized the idea could be turned around: he could use the wavelengths of x rays emitted from heavier elements to determine how many protons they had in their nuclei. His work set the stage for a new interpretation of the periodic table.
Moseley’s results led to the conclusion that the order of elements in the periodic chart was based on some fundamental principle of atomic structure. Because of Moseley’s work, scientists were convinced that the periodic nature of the properties of elements is due to differences in the numbers of subatomic particles. As each succeeding element is added across a row on the periodic chart, one proton and one electron are added. The number of neutrons added is unpredictable but can be determined from the total weight of the atom. However, since Mendeleev’s time the character of the periodic table has remained basically unchanged, providing testament to the power of his original insight.
After Moseley’s work, the idea that the periodic patterns in chemical reactivity might actually be due to the number of electrons and protons in atoms intrigued many chemists. Among the most notable was American physical chemist Gilbert Newton Lewis (1875-1946) of the University of California at Berkeley. Lewis explored the relationship between the number of electrons in an atom and its chemical properties, the kinds of substances formed when elements reacted together to form compounds, and the ratios of atoms in the formulas for these compounds. Lewis concluded that chemical properties change gradually from metallic to nonmetallic until a certain stable number of electrons is reached.
An atom with this stable set of electrons is a very unreactive species. However, if one more electron is added to this stable set of electrons, the properties and chemical reactivities of this new atom change dramatically: the element is again metallic, with the properties like elements of Group 1. Properties of subsequent elements change gradually until the next stable set of electrons is reached and another very unreactive element completes the row.
Layout of the periodic table
The first step in being able to use the information contained in the periodic table is to understand how it is arranged. Most periodic tables are similar to one another but to lessen confusion the periodic table shown in Figure 1 will be used. One of the first things that stands out is that the table is composed of metals, nonmetals, and metalloids.
The metallic elements are familiar to people all through everyday living. From experience humans know that metals are shiny, conduct heat and electricity very well (think about electrical wires and pots and
pans), can be formed into many different shapes (in other words, they are malleable), and can be drawn into wires (are ductile). The only metal that is not a solid at room temperature is mercury, which exists as a liquid and is often used in thermometers. The non-metal elements familiar to people include the atmospheric gases nitrogen and oxygen (O). Other important nonmetals, especially for the maintenance of life, are carbon (C), hydrogen (H), sulfur (S), and phosphorus (P). Most nonmetals are either gases or solids at room temperature and have properties opposite those of the metals.
The placement of metals and nonmetals in the table, it should be noticed, is not random. The non-metals all occur on the right hand side of the table, while the metals occur on the left hand side. Moving across the table from metals to nonmetals the metalloids (or semimetals) are encountered, which include boron (B), silicon (Si), germanium, arsenic (As), antimony (Sb), tellurium, and astatine (At). The properties of metalloids fall in between those of metals and nonmetals. A familiar metalloid is silicon, the major material of which computer memory chips are made.
The periods, or horizontal rows, of the table are numbered on the left hand side from 1 to 7. The first period contains two elements, hydrogen and helium (He), the second period and third period each contain eight elements, while the fourth and fifth periods each contain 18 elements.
The numbering of groups (the vertical columns, also known as families) follows two different conventions, both of which should be familiar. In the system commonly used in North America, Roman numerals and letters are used to denote the various groups. The alternate system, devised by the International Union of Pure and Applied Chemistry Convention (IUPAC— the same group responsible for certifying atomic masses and element names) in 1985, numbers the Groups from 1 through 18. The IUPAC system is the system to which most countries are turning and this is the system used in this text. The alternative system numbering will be shown in parenthesis when applicable.
In both systems the various groups in the periodic table are placed into families that consist of groups of related elements. These families are given the same name in each system. Groups 1 and 2 (IA, IIA) are called the main group metals. Group 1 individually, is referred to as the alkali metals while Group 2 is called the alkaline earth metals. The Group 1 and 2 metals are both very reactive and readily form positive charged atoms (called cations) by losing electrons. Group 1 metals lose one electron to become +1 cations and Group 2 metals can lose two electrons to become +2 cations.
Groups 3 through 12 (refer to table to see alternate numberings) are referred to as the transition metals. The transition metal family, unlike the main Group 1 and 2 metals, form cations of differing charge (from +1 to +3). Many transition metal compounds are colored. Groups 13 through 18 (lllA through VlllA) are called the main group nonmetals. The inner transition metal family is comprised of two series called the lanthanides and actinides, neither of which are numbered in either system.
As briefly explained before, each box, along with the symbol for the element, represents an individual element. Each element is characterized by a unique atomic number (the number that appears above the elemental symbol), which denotes the number of protons in the nucleus of that atom. One can see that the number of protons determines the element. If an atom has six protons in its nucleus it is a carbon atom, while 34 protons determine a selenium (Se) atom. Protons each carry a charge of +1, while electrons carry a charge of -1. Therefore, neutral elements must have equal numbers of protons and electrons.
Also contained in each box is a number written below each elemental symbol. This number is the atomic mass number, the average atomic mass for a given element. It is an average because not all atoms of a given element have the same mass. While all atoms of the same element must have the same number of protons, as mentioned above, they can differ in the number of neutrons they contain in the nucleus. Neutrons, as the name implies, are neutral particles found in the nucleus. These particles help to stabilize the atom by separating the positively charged protons in the nucleus. Some elements have only one possible combination of protons and neutrons, such as sodium. All sodium (Na) atoms consist of 11 protons and 12 neutrons. Atoms containing the same number of protons but different numbers of neutrons are referred to as isotopes. For example, there are two isotopes of carbon found in nature. Carbon-12 has six protons and six neutrons in its nucleus and a natural abundance of 98.889%, while carbon-13 has six protons and seven neutrons with a natural abundance of 1.111%. By averaging the atomic mass of each isotope of carbon, the average atomic mass of 12.01 atomic mass units, or amu (1 amu = 1.66 x 10-27 kilograms), is calculated and appears under carbon in the periodic table. It is important to note that isotopes of the same element are in most instances chemically indistinguishable.
Electronic structure
At this point it should be clear what makes one element different from another (differing numbers of protons), but what makes them similar? What allowed Mendeleev to arrange the elements into a periodic table whereby elements with similar chemistry were placed one under the other?
The last piece of atomic sub-structure needed to fully explain the arrangement of the elements of the periodic table is the electron configuration. It is the arrangement of the electrons around the nucleus that determines the degree and type of reactivity an element will exhibit. At the time that Mendeleev assembled the periodic table, electrons as well as the sub-atomic structure of the atom were yet to be discovered. Their discovery revealed the underlying principles upon which the periodic table is based.
Elements that appear in the same group have the same valence shell electron configuration. Electrons are found in shells, or energy levels, around the nucleus of the atom. The farther the shell is from the nucleus, the higher the energy of its electrons. The valence shell is the outermost shell of an atom. The number of shells in an atom of an element can be determined by noting the number of the period in which that element is found. For example, potassium (K) is in Period 4, which means it has four electron shells. The valence shell is the fourth shell, and the farthest from the nucleus. Sodium is in Period 3, which means it has three shells. The valence shell is the third shell, which is closer to the nucleus than the fourth. By this reasoning, it is easy to see that the valence electrons in potassium must be of higher energy than the valence electrons in sodium. It is the electrons in the valence shell that are involved in chemical reactions. Electrons below the valence shell are considered core electrons and are not important when determining reactivity. To be able to fully use the table, subshells, which help describe the locations of an atom’s electrons, need to be briefly explained.
As just discussed, the outermost electron shell of an atom is the same as the period in which an element is found. The shell can be thought of as a street name in the address of the electrons. Each of the families of elements belongs to a particular subshell, which can be thought of as the house number in the address of the electrons. The valence electrons of the main group metals (Groups 1 and 2) are in the s subshell. The valence electrons of the main group nonmetals (Groups 13 through 18) are in the p subshells, the transition metal valence electrons are in the d sub-shells, while the inner transition metal valence electrons are in the f subshells. The energy of the subshells increases, within the same shell, from s, then p, then d, and finally f.
Each subshell contains orbitals, which are like the rooms of the house at the particular address of the electrons. Each orbital holds a maximum of two electrons. There is only one orbital in each s subshell, three orbitals in each p subshell, five orbitals in each d subshell, and seven orbitals in each f subshell. Therefore, the s subshell can hold two electrons, the p subshell can hold six electrons, the d subshell can hold 10 electrons, and the f subshell can hold 14 electrons.
To illustrate how these electronic properties are relevant to the periodic table the first three elements of Group 16, oxygen, sulfur, and selenium will be examined. Each of these elements has a valence shell electron configuration of s2p4, which means there are two electrons in the s subshell and four electrons in the p subshell. Although they are each in different periods, their electronic structure is the same and has similar chemistry. For example, the compound water has the formula H2O. Likewise there are the compounds H2S and H2Se. In a similar fashion, if told that the Group 15 element nitrogen, with the valence shell electronic configuration s2p3, forms the compound NH3 (called ammonia), can one infer the formula of the compound that forms between phosphorus and hydrogen? By analogy to NH3 one can expect the compound to have the formula PH3. It is by this same type of reasoning that Mendeleev predicted the existence of the unknown elements.
Group 18 in the periodic table, called the noble gases, are all very unreactive elements. They do not easily combine, if at all, with other elements. This indicates that there is some special stability to the electron configuration s2p6 that the noble gases possess. When an element has the full shell configuration s2p6 it is referred to as having an octet, or eight valence electrons. Much of the reactivity of the elements can be described as an attempt to achieve an octet. Since the noble gases naturally have an octet configuration, they do not need to react with other atoms to achieve this stable structure.
In the ionic compound sodium chloride (NaCl) one finds a positively charged sodium atom (Na+)
and a negatively charged chlorine atom (Cl-). If the valence shell electron configuration of each ion is examined, one can see that a chlorine atom, by gaining an extra electron, goes from a s2p5 (Group 17) configuration to the stable s2p6 (Group 18) configuration of the chloride ion. The chloride ion is referred to as being isoelectronic (having the same electronic configuration) with argon (Ar). The sodium atom (s1) can lose an electron to become a sodium cation with the stable s2p6 configuration, making it isoelectronic with neon (Ne).
Other trends
There are general reactivity trends on the periodic table that are useful to know. Metals and nonmetals usually combine to form ionic compounds with the metal giving up an electron to become positively charged and the nonmetal element gaining an electron to become negatively charged. Nonmetals usually combine with one another to form covalent bonds in which the electrons are not fully transferred but are shared between the two elements. Examples of this are molecular oxygen O2, molecular chlorine Cl2, ammonia NH3, and carbon dioxide CO2.
The degree of metallic character of an element can be estimated by that element’s location in the periodic table. Metallic character decreases moving from left to right across a period. This is clearly demonstrated in each of the first six periods where each period starts off with metallic elements, but ends with nonmetallic elements. Metallic character is also found to increase moving down a group. This trend is most clearly demonstrated by Groups 14-16, although it is true for the other groups as well, for example, cesium (Cs) is more metallic than sodium. Each of the Groups 14-17 begins with a nonmetal followed closely by a metalloid and eventually a metallic element.
Another trend found in the periodic table is size. The atomic radii (the scientific term for the size of an atom) of the elements increases going down a group, while it decreases going across a period. See Figure 2.
In the periodic table, the last naturally occurring element is element 92, uranium (U). Uranium is a radioactive element. Radioactive elements are unstable and break down to form lighter elements and in the process give off energy. All of the elements that occur past uranium are artificially made, and are referred to collectively as the transuranium elements.
Recent and future research
While the general form of the periodic table has withstood the test of time and should change very little in the future, alterations of the periodic table have been and continue to be made. One area that could see minor changes is the atomic mass. In the future, more accurate methods of measuring the mass of atoms may be invented. The magnitude of these changes, however, would be exceedingly small.
The largest area of change in the periodic table will come from the artificially made creation of new chemical elements. Every element past uranium in the periodic table has been made by scientists in high-energy particle accelerators. The first transuranium element made was element 93, discovered by American scientist Edwin Mattison Macmillan (1907—1991) and P.H. Abelam at the University of California at Berkeley in 1940. The two discoverers of this element named it neptunium (Np).
The discovery of elements 95, americium (Am), and 96, curium (Cm), caused a dilemma. It was thought that these new elements should be placed after element 89, actinium (Ac) in the d-block transition metal family. American chemist Glenn Theodore Seaborg (1912—1999), Nobel Prize winner in 1951 for the discovery of plutonium (Pu) as well as nine other transuranium elements, felt that they should be placed under the lanthanides in a new group as part of the inner transition metal family. Further experimentation showed that they did belong in the inner transition metal family. The discovery of elements 104 through 111, which belong in the transition metals family, proved that the proposed groupings were correct.
Unlike most of the naturally occurring elements, which can be handled and studied, the transuranium elements are all radioactive and break down incredibly fast. The synthesis and detection of transuranium elements takes great technical expertise. In addition, the experimental machinery needed to do this work is extremely expensive as well as complicated, therefore only a few research centers in the world are involved in this area of study.
The transuranium elements are synthesized by colliding accelerated charged particles with heavy atoms (i.e., curium and lead). In certain collisions, the nuclei of the accelerated charged particles and the stationary heavy atoms will fuse to produce a new transuranium element. The lifetimes of these new elements are so
KEY TERMS
Anion— A negatively charged ion (i.e., Cl-).
Atomic mass— The mass of an atom relative to carbon-12 (which has a mass of exactly 12 atomic mass units); also the mass, in grams, of an element that contains one mole of atoms.
Atomic number— The number of protons in the nucleus of an atom.
Cation— A positive ion (i.e., Na+).
Covalent bond— A chemical bond formed when two atoms share a pair of electrons with each other.
Electron— A negatively charged particle, ordinarily occurring as part of an atom. The atom’s electrons form a sort of cloud about the nucleus.
Electron configuration— The arrangement of electrons in the occupied electron energy levels or sub-levels of an atom.
Element— A pure substance that can not be changed chemically into a simpler substance.
Family— A set of groups characterized by the same subshell.
Group— A vertical column of the periodic table that contains elements possessing the same electronic configuration.
Ionic bond— The attractive forces between positive and negative ions that exist when electrons have been transferred from one atom to another.
Isotopes— Two molecules in which the number of atoms and the types of atoms are identical, but their arrangement in space is different, resulting in different chemical and physical properties.
Neutron— A subatomic particle with no electric charge.
Nucleus— Small core at the center of atoms that contain the protons and neutrons.
Octet (noble gas configuration)— The stable electron configuration found with Group 18 elements, also referred to as the closed shell configuration.
Period— Horizontal rows of the periodic table.
Proton— Subatomic particle of +1 charge.
Shell— Energy level within an atom. The period of an element determines the shell number.
Subshell— Further energy levels found within a given shell. Elements in the same family share the same subshell.
Transuranium— Term given to all the manmade elements of greater atomic number than 92.
short they often break down into other elements within fractions of a second and are detected only by their breakdown products, referred to as daughter elements.
Recently, several new artificially made super-heavy elements have been discovered. These include elements 110 and 111, both of which were made in late 1994 by an international team of scientists. These scientists performed this research at GSI, a research center for Heavy Ion Research Laboratory (HIRL) in Darmstadt, Germany. Element 110 was made by colliding nickel atoms with an isotope of lead. Researchers in Russia have plans to make a different isotope of element 110 by colliding sulfur atoms with plutonium atoms.
Element 112 was discovered in 1996 at HIRL, while element 113 was discovered in 2003 at Lawrence Livermore National Laboratory (LLNL) in California and the Joint Institute for Nuclear Research (JINR) in Dubna, Russia. Element 114 was discovered in 1998 at JINR, while element 115 was discovered in 2003 at LLNL and JINR. Element 116 was discovered in 2000 at JINR. Elements 112 through 116 have yet to be independently confirmed. Other super-heavy elements, which have been predicted to exist (such as elements 117 and 118), have yet to be completely created in the laboratory, although research continues into the creation of these elements.
Based on theoretical calculations some researchers believe that not all transuranium elements will be so unstable. Already, different isotopes of element 106, seaborgium (Sg) have been made that are stable for up 33 times longer than the original isotope discovered in 1974 (even at 33 times longer lifetime it only is stable for a maximum of thirty seconds). It is theorized that some isotopes of the yet to be made transuranium elements should be stable for very long periods of time, allowing them to be studied chemically. Many exciting discoveries remain to be uncovered concerning the creation of new elements, and with the periodic table as a guide, their place is already awaiting them.
Names of the elements
The naming and symbol of the elements in the periodic table is an interesting story itself. Many of the element symbols are derived from the elemental name such hydrogen (H), oxygen (O), chlorine (Cl), and calcium (Ca). Other element symbols seem to bear no relationship to their name such as sodium (Na), tin (Sn), and lead (Pb). These elemental symbols all derive from the Latin name of the element: natrium, stan-num, and plumbum. Many of the elements have been named by their discoverer.
The element phosphorus was named by its discoverer for the property that it glows when exposed to air. Phosphorous in Greek means: I bear light. From the names of the elements such as francium (Fr), americium, europium (Eu), berkelium (Bk), and californium (Cf), it is clear that geographic locations were used to name them. Still other elements have been named to honor people. In this category falls element 101, mendelevium (Md), named to honor the discoverer of the periodic table. Others in this category include einstein-ium (Es) and nobelium (No), named after German— American physicist Albert Einstein (1879—1955) and Swedish inventor Alfred Nobel (1833—1896).
At this time, to name an element a researcher or team of researchers must be certified by IUPAC as the discoverers of that element, at which time they are free to name the compound. The elements 104 through 109 were subject to a naming controversy. The originally proposed names of these elements by IUPAC were, in order, dubnium, joliotium, rutherfordium, bohrium, hahnium, and meiterium. The names that appear on the current periodic table are, in order, rutherfordium (Rf), dubnium (Db), seaborgium (Sg), bohrium (Bh), hassium (Hs), and meitnerium (Mt).
A particular controversy among these elements involved element 106 that researchers at Berkeley were credited with discovering by IUPAC. Following historical convention, the Berkeley researchers were free to name the element. They chose to name it seaborgium, after Glenn T. Seaborg who contributed to the element’s discovery. IUPAC ignored the recommendations of the discoverers and suggested the name rutherfordium for element 106. A vote of the IUPAC Council in August 1995 resolved the issue, and now element 104 is called rutherfordium and element 106 is called seaborgium.
As a final testament to the great respect with which the periodic table is held, it is instructive to hear Glenn T. Seaborg talk about the significance of having his name assigned to element 106: “A thousand years from now, seaborgium will still be in the periodic table, whereas the twentieth-century Nobel Prize-winners will seem a very small part of history. . . . This honor will last as long as civilization.”
See also Atomic weight; Element, chemical; Subatomic particles.
Resources
BOOKS
Brock, William Hodson. Norton History of Chemistry. New York: Norton, 2000.
Emsley, John. Nature’s Building Blocks: An A-Z Guide to the Elements. Oxford, UK: Oxford University Press, 2003.
Hoffmann, Roald, and Torrence, Vivian. Chemistry Imagined Reflections on Science. Washington: Smithsonian Institutional Press, 1993.
Roberts, Royston M. Serendipity: Accidental Discoveries in Science. New York: John Wiley & Sons Inc., 1989.
Siekierski, Slawomir. Concise Chemistry of the Elements.Chichester, UK: Horwood Publishing, 2002.
OTHER
Naeye, Robert. “An Island of Stability.” Discover (August 1994): 15.
Michael G. Roepel