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Periodic Table (Predicting the Structure and Properties of the Elements)

Periodic table (Predicting the structure and properties of the elements)

An element is defined by the number of protons in the nucleus of its atoms, but its chemical reactivity is determined by the number of electrons in its outer shella property fundamental to the organization of the periodic table of the elements.

In the second half of the nineteenth century, data from laboratories in France, England, Germany, and Italy were assembled into a pamphlet by Stanislao Cannizzaro (18261910), 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 1860 international meeting held in Karlsruhe, Germany, that was organized to discuss new ideas about the theory of atoms. When Russian chemist and physicist Dmitri Mendeleyev (18341907) 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 intervalsthey displayed periodicity. Mendeleyev used the periodic repetition of chemical and physical properties to construct a chart much like the Periodic Table we currently use.

Early in the twentieth century work initiated by Joseph John Thomson (18561940) in England led to the discovery of the electron and, later, the proton. In 1932, James Chadwick (18911974), also in England, 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.

In 1913, a third British scientist, Henry G. J. Moseley (18871915), 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. As a result 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 .

When Mendeleyev placed elements in his Periodic Table, he had all elements arranged in order of increasing relative atomic weight. However, in the modern Periodic Table, the elements are placed in order of the number of protons in the nucleus. As atomic weight determinations became more precise, discrepancies were found. The first case of a heavier element preceding a lighter one in the modern Periodic Table occurs for cobalt and nickel (58.93 and 58.69, respectively). In Mendeleyev's time, both atomic weights had been determined to be 59. Mendeleyev grouped both elements together, along with iron and copper. From the work of Moseley and others, the number of protons in the nuclei of the elements cobalt and nickel had been found to be 27 and 28, respectively. Therefore, the order of these elements, and the reason for their similar behavior with respect to other members of their chemical families, arises because nickel has one more proton and one more electron than cobalt.

Whereas Mendeleyev based his order of elements on mass and chemical and physical properties, the arrangement of the table now arises from the numbers of subatomic particles in the atoms of each element. The stage was now set for examining how the number of subatomic particles affects the chemistry of the elements. The role of the electrons in determining chemical and physical properties was obscure early in the twentieth century, but that would soon change with the pivotal work of American chemist Gilbert N. Lewis (18751946).

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 Lewis, then at 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. But 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.

A stable number of electrons is defined as the number of electrons found in an unreactive or "noble" gas. Lewis suggested electrons occupied specific areas around the atom, called shells. The noble gas atoms have a complete octet of electrons in the outermost shell.

The observation that each element starting a new row has just one electron in a new shell opens the door to relating chemical properties to the number of electrons in a shell. Mendeleyev put elements together in a family because they had similar reactivities and properties; Lewis proposed that elements have similar properties because they have the same number of electrons in their outer shells.

Many observations of the chemical behavior of elements are consistent with this idea: the number of electrons in the outer shell of an atom (the valence electrons) determines the chemical properties of an element. Lewis extended his ideas about the importance of the number of valence electrons from the properties of elements to the bonding of atoms together to form compounds. He proposed that atoms bond with each other either by sharing electrons to form covalent bonds or by transferring electrons from one atom to another to form ionic bonds. Each atom forms stable compounds with other atoms when all atoms achieve complete shells. An atom can achieve a complete shell by sharing electrons, by giving them away completely to another atom, or by accepting electrons from another atom.

Many important compounds are formed from the elements in rows two and three in the Periodic Table. Lewis predicted these elements would form compounds in which the number of electrons about each atom would be a full shell, like the noble gases. The noble gases of rows two and three, neon or argon, each have eight electrons in the outermost valence shell. Thus, Lewis's rule has become known as the octet rule and simply states that there should be eight electrons in the outer shell of an atom in a compound. An important exception to this is hydrogen for which a full shell consists of only two electrons.

The octet rule is followed in so many compounds it is a useful guide. However, it is not a fundamental law of chemistry. Many exceptions are known, but the octet rule is a good starting point for learning how chemists view compounds and how the periodic chart can be used to make predictions about the likely existence, formulas and reactivities of chemical substances.

Elements in a vertical column of the Periodic Table typically have many properties in common. After all, Mendeleyev used similarities in properties to construct a periodic table in the first place. Because they show common characteristics, elements in a column are known as a family. Sometimes a family had one very important characteristic many chemists knew about: that characteristic became the family name. Four important chemical family names of elements still widely used are the alkali metals , the alkaline earths, the halogens, and the noble gases. The alkali metals are the elements in Group 1, excluding hydrogen, which is a special case. These elementslithium, sodium, potassium, rubidium, cesium and franciumall react with water to give solutions that change the color of a vegetable dye from red to blue. These solutions were said to be highly alkaline or basic; hence the name alkali metals was given to these elements.

The elements of Group 2 are also metals. They combine with oxygen to form oxides, formerly called "earths," and these oxides produce alkaline solutions when they are dissolved in water. Hence, the elements are called alkaline earths.

The name for Group 17, the halogens, means salt former because these elements all react with metals to form salts.

The name of Group 18, the noble gases, has changed several times. These elements have been known as the rare gases, but some of them are not especially rare. In fact, argon is the third most prevalent gas in the atmosphere, making up nearly 1% of it. Helium is the second most abundant element in the universeonly hydrogen is more abundant. Another name used for the Group 18 family is the inert gases. However, Neil Bartlett, while at the University of British Columbia in Vancouver, Canada, showed over 30 years ago that several of these gases could form well-defined compounds. The members of Group 18 are now known as noble gases. They do not generally react with the common elements but do on occasion, especially if the common element is as reactive as fluorine.

Knowing the chemistry of four families of the periodic tablegroups 1, 2, 17, and 18, the alkali metals, the alkaline earths, the halogens and the noble gasesenables chemists to divide the elements in the Periodic Chart into other general categories: metals and nonmetals. Metals are hard but ductile substances that conduct electricity . Groups 1 (excluding hydrogen) and 2 are families of metallic elements. Groups 17 and 18 contain elements with very different properties perhaps best described by what they are notthey are not metals, and hence are called nonmetals. Between Groups 1 and 2, and Groups 17 and 18 is a dividing line between these two types of elements. Most periodic charts have a heavy line cutting between aluminum and silicon and descending downward and to the right in a stair-step fashion. Elements to the left of the line are metallic; those to the right, nonmetallic. The boundary is somewhat fuzzy, however, because the properties of elements change gradually as one moves across and down the chart, and some of the elements touching that border have a blend of characteristics of metals and nonmetals; they are frequently called semi-metals or metalloids.

The elements in the center region of the table, consisting of dozens of metallic elements in Groups 312, including the lanthanide and actinide elements, are called the transition elements or transition metals. The other elements, Groups 1,2, and 1318, are called the representative elements.

There is a correlation among the representative elements between the number of valence electrons in an atom and the tendency of the element to act as a metal, nonmetal, or metalloid. Among the representative elements, the metals are located at the left and have few valence electrons. The nonmetals are at the right and have nearly a full shell of electrons. The metalloids have an intermediate number of valence electrons.

The structure and bonding of a compound determine its chemical and physical properties. Lewis's idea of stable, filled electron shells can be used to predict what atom is bonded to what other atom in a molecule. In many cases, Lewis's octet rule is followed by taking one or more electrons from one atom to form a cation and donating the electron or electrons to another atom to form an anion. Metallic elements on the far left of the Periodic Table can lose electrons and elements on the far right can readily accept electrons. When these elements combine, ionic bonds result. An example of an ionic compound is sodium chloride. The sodium cation, Na+, forms an ionic bond with chloride anion, Cl.

In covalent bonds, electrons are shared between atoms. Lewis defined a covalent bond as a union between two atoms resulting from the sharing of two electrons. Thus, a covalent bond must be considered a pair of electrons shared by two atoms. Elemental bromine, Br2, is an example of a covalent compound. Each bromine atom has seven electrons in its outer shell and requires one electron to achieve a noble gas configuration. Each can pick up the needed electron by sharing one with the other bromine atom.

Water, the solvent of life and an important agent in many geochemical processes, is a compound formed by the combination of atoms of two nonmetallic elements, hydrogen and oxygen. Each hydrogen atom requires just one electron to fill its shell because the first shell (the number of electrons of the noble gas helium) holds only two electrons. Oxygen lacks two electrons compared with neon, the nearest noble gas. If each hydrogen can obtain one electron by sharing electrons from the oxygen atom and the oxygen atom can share one electron from each of the two hydrogen atoms, every atom will have a full shell of electrons, and two covalent bonds will be formed as a result of sharing two pairs of electrons.

One of the most important properties of an element that can be used to predict bonding characteristics is whether the element is metallic or nonmetallic.

Pure metals are typically shiny and malleable. Chemists have found metals also have common chemical properties. Metals combine in similar ways with other elements and form compounds with common characteristics. Metals combine with nonmetals to form salts. In salts, the metals tend to be cations. Salts conduct electricity well when melted or when dissolved in water or some other solvents but not when they are solid.

Most pure metals, when freshly cut to expose a new surface, are lustrous, but most lose this luster quickly by combining with oxygen, carbon dioxide, or hydrogen sulfide to form oxides, carbonates or sulfides. Only a few metals such as gold, silver, and copper are found pure in nature, uncombined with other elements.

Nonmetals in their elemental form are usually gases or solids. A few are shiny solids, but instead of being metallic gray they are typically black (boron, carbon as graphite ), colorless (carbon as diamond ), or highly colored (violet iodine, yellow sulfur). At room temperature , only one of them is a liquid (bromine).

Nonmetallic elements combine with metallic elements to form salts. In salts, the nonmetallic elements tend to be anions. Non-metals accept electrons in forming anions while metals donate electrons to form cations. This reflects a periodic property of elements: as one moves from left to right across a row on the periodic chart, on the left are the atoms of metals which tend to give up electrons relatively easily and on the right side are nonmetals which do not readily give up electrons in forming chemical bonds . At the start of the next row, the trend is repeated. This periodic property is referred to as electronegativity. The more readily atoms accept electrons in forming a bond, the higher their electronegativity. Metals are characterized by low electronegativities; nonmetals, by high electronegativity. Electronegativity increases across a row on the periodic chart.

Nonmetallic elements combine with each other to form compounds. Although some nonmetallic elements form solutions when mixed with other nonmetallic elements, most react with other nonmetals to form new substances. For example, at the high temperatures and pressures of an internal combustion engine, nitrogen and oxygen gases from the atmosphere react to form nitrogen oxides such as nitric oxide, NO, and nitrogen dioxide, NO2. Nonmetallic elements form covalent bonds with each other by sharing electron pairs. This tendency to bond by sharing electrons reflects the periodic trend described above: elements on the right side of the periodic chart do not give up electrons easily when forming bonds; their electronegativity is high. They tend to either accept electrons from metals to form salts or share electrons with other nonmetals to form covalent compounds.

Metalloids typically show physical characteristics (e.g., electrical conductivity) intermediate between the metals and nonmetals. Metalloids typically act more like nonmetals than metals in their chemistry. They more often combine with nonmetals to form covalent compounds rather than salts, but they can do both. This reflects their intermediate position on the Periodic Table. They can form alloys with metals and with the other metalloids. Semiconductors are typically made from combinations of two metalloids. The minor constituent, for example germanium, is said to be "doped" into the major constituent, which is often silicon.

The boundaries between metals, nonmetals, and metalloids are arbitrary. The changes in properties as one moves from element to element on the chart are gradual.

Earth's atmosphere contains slightly more than 20% oxygen. Because oxygen is quite reactive, most elements can be found in nature as oxides. The alkali metals (Group 1) and alkaline earths (Group 2) were so named because the metallic oxides formed when the metals reacted with oxygen produced basic solutions when dissolved in water. Metallic oxides are known as basic anhydrides (anhydrous, meaning without water), because basic solutions are formed when they are added to water.

Nonmetallic elements combine with oxygen to form oxides, many of which, such as carbon dioxide, sulfur dioxide and nitrogen dioxide, are gases. When oxides of nonmetallic elements are dissolved in water, they tend to form acidic solutions or neutral solutions. Nonmetal oxides that form acidic solutions when dissolved in water are called acid anhydrides.

Transition metals react with oxygen to form a wide variety of oxides, some of which are basic and some acidic. A few transition metals are relatively unreactive and may be found in nature as pure elements.

See also Atmospheric chemistry; Atomic mass and weight; Atomic number; Atomic theory; Atoms; Chemical bonds and physical properties; Chemical elements

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Periodic Table

Periodic Table


The Periodic Table places the symbols of chemical elements, sequenced by atomic number , in rows and columns that align similar properties.

Antiquity through the Renaissance

A few thousand years ago, primitive chemistry focused mostly on converting one substance into another. The word "chemistry" itself is arguably traced to the name of a region of ancient Egypt where such transformation attempts were practiced. Over the centuries, philosophers tried to come to terms with the growing variety of known substances. They postulated the role of fundamental entities that could not be broken down further but formed simple materials when combined. By the time of ancient Greece, Democritus, Leucippus, and Empedocles expounded the nature of matter in terms of constituent elements, the simple substancesearth, air, fire, and waterof which all materials were compounded. The term "atom" first appears in this context.

A millennium or so later, Arab civilizations made great strides in laboratory techniques. Subsequently, during the Renaissance period, these techniques were adopted in trying to transform one element into another, most notably into gold from less costly substances like lead. This gave the Arabic term "alchemy" its modern mystical connotation.

Post-Renaissance

By the mid-1700s about twenty elements were known. Science was beginning to get more sophisticated as measurements and instrumentation improved rapidly and theories based on observation grew more advanced and more compelling. Chemists, however, continued to anguish over the inability to easily categorize the elements.

What was likely the first attempt at sorting the elements was a table of simple substances, prepared in 1772 by French chemist Louis-Bernard Guyton de Morveau. French chemist Antoine-Laurent Lavoisier was most influential in developing an experimental approach, which is acknowledged to have laid the foundation for modern chemistry. In 1789, Lavoisier published a list of pure substances that included the known elements but also some compounds and light and heat. By the early 1800s, following the introduction of English chemist John Dalton's atomic theory and the concept of atomic weights, the number of known elements had grown. Although properties were carefully measured, confusion held sway when it came to agreeing on the composition of compound substances and the related assignment of atomic weights.

In 1829 German chemist Johann Döbereiner noted that there were triplets of elements in which the central species' properties were almost exactly midway between the outer two. The first example of such a triplet included chlorine, bromine, and iodine. Properties such as atomic weights, color, and reactivity followed this "law of triads" for several such groupings, but not for the entire collection of known elements.

In 1860 Italian chemist Stanislao Cannizzaro presented analyses at an international chemistry meeting that, when merged with previously ridiculed hypotheses by fellow Italian Amedeo Avogadro, yielded unambiguous atomic weights. These eliminated most of the disharmony among property determinations. In attendance were German chemist Lothar Meyer and Russian chemist Dimitri Mendeleev, both of whom were inspired to give the presentation further thought.

Industrial Age

In 1862 French geologist Alexandre-Émile Beguyer de Chancourtois arranged the elements in order of increasing atomic weights, wrapping the series around a cylinder in a helical display. He noted that elements with similar properties lined up, one over the other. His idea was obscured by its publication in a nonchemistry journal, the inclusion of compounds and alloys in the discussion, and the publisher's decision not to include an essential diagram.

JOHN NEWLANDS (18371898)

John Newlands compared elements to musical notes with his law of octaves. As on a scale, every eighth element would share similar properties when arranged by increasing atomic weight. Newlands did not account for exceptions, however, and it was only upon establishment of the Periodic Table that his theory gained credibility.

Valerie Borek

A few years later, British chemist John Newlands also arranged the elements in order of increasing atomic weights. He was apparently the first to assign hydrogen the weight of "1." Newlands noted that properties repeated when the sequence was broken into periods of seven and referred to his system as the "law of octaves."

During the mid-1860s Meyer took the newly established atomic weights of many elements and arranged them into families that bore similarities in properties, including the ability of an atom to combine with other atoms (valency).

In 1869 Mendeleev presented his table of the elements (sixty-three by now) arrayed in periods of seven for the lighter elements and opening up to seventeen for the heavier elements. Furthermore, Mendeleev had the foresight and confidence to break the atomic weight sequence by occasionally forcing elements out of order so as to fall in an appropriate location as determined by their properties. He left gaps in the arrangement at several places and used implied trends to predict characteristics of undiscovered elements needed to complete the table.

Mendeleev's Periodic Table was not well received at first, but was shortly helped by the discovery of the element gallium, which filled such a gap and had nearly exactly the atomic weight, density, and valency predicted. Other affirmations followed. It is for these reasons that Mendeleev is given most of the credit for the invention of the Periodic Table.

In 1892 Scottish chemist William Ramsay discovered two more elements, argon and helium. These unreactive gases did not fit into the Periodic Table. In short order, Ramsay also discovered three more unreactive gases. These gases represented a new family of elements that had to be inserted as an eighth main column in Mendeleev's table.

The Modern Periodic Table

As the twentieth century approached, elements of similar properties were arranged in eight main vertical columns referred to as chemical families. The first such family, or Group I, is collectively termed the "alkali metals," commencing with lithium. The next column, Group II, is designated the "alkaline earths," commencing with beryllium. Groups III through V are commonly referred to as the boron, carbon, and nitrogen families, respectively. The next group, the oxygen family, is technically called the "chalcogens." Group VII, the "halogens ," begins with fluorine. Finally, the elements of Group VIII, starting with helium, are called the "noble gases." Because of their relative unreactivity, they had once also been termed inert gases, a label no longer acceptable.

In the absence of any understanding as to why the periodic arrangement appeared as it did, or whether or not there were yet more surprises, the science of chemistry remained incomplete, although very important and practical. At almost exactly this time, just before the start of the twentieth century, three findings were announced that changed the course of science: x rays were discovered by German physicist Wilhelm Röntgen in 1895, radioactivity by French physicist Antoine-Henri Becquerel in 1896, and the electron by British physicist Joseph John Thomson in 1897.

First Model

What soon emerged was a nuclear model of the atom, first proposed by New Zealand-born physicist Ernest Rutherford. In this view, an element's identity was determined by its atomic number, the amount of positive charge in the very small core nucleus that also contained almost all of the atom's mass. The light electrons were held in orbits by electrostatic attraction to the positive core.

Rutherford's view was extended by Danish physicist Niels Bohr in 1913. Bohr modeled that electrons moved in fixed orbits around the nucleus, much as planets orbit the Sun. Furthermore, not only were the locations of these orbits fixed, but so were the speeds of the electrons in each orbit and the number of electrons that could be accommodated in each orbit, a description called the electron configuration. By explaining the quota of electrons allowed in each fixed orbit, Bohr resorted to a new physical idea called quantization. As a consequence, Bohr was able to reproduce the Periodic Table, adding one electron at a time as one stepped to the next higher element. Bohr argued that orbits of increasing radius could accommodate up to a maximum number of electron numbers that, when reached, corresponded to observed horizontal periods of two, eight, eight, eighteen, eighteen, and thirty-two. Bohr acknowledged the unattractiveness of this approach in that it was merely mimicking an observed pattern rather than addressing the underlying science.

Modern Theory

The mid-1920s witnessed a necessary breakthrough. The revolutionary wave concept of matter was incorporated into a mathematical framework, a new quantum theory, that explained all the properties of a bound electron: its energy, the description of where it could be found, and configuration restraints.

An electron could have only certain energies determined by the value of an integer (a whole number), traditionally symbolized by n with values 1, 2, 3, and so on. Electron energy with n = 1 is the lowest possible, n = 2 being the next lowest, and so on. The region of space where the electron might be foundcalled an orbital because it replaced Bohr's planetary fixed orbit ideacould be characterized by its size, shape, and orientation (how the shape might be tilted). For each n, there was a determined set of shapes and orientations with letters used to indicate the shapes. For n = 1, only a spherical shape is allowed, symbolized by s ; since a sphere has no orientation, that is the only n = 1 orbital. It is abbreviated as 1s. For n = 2, there are larger orbitals: another s, the 2s, and also dumbbell-shaped orbitals with opposing lobes. These are symbolized as p orbitals and have three possible orientations for the 2p and all other p orbitals. By the time n = 3 is considered, there is a third shape, d, with five orientations. For n = 4 there is a fourth shape, f, with seven orientations in addition to the 4s, 4p and 4d. The sequence of filling follows a relatively simple pattern shown by arrows in Figure 1.

Very early in the development of modern quantum mechanics, German physicist Wolfgang Pauli realized that each of the substates characterized by n, shape, and orientation was permitted to have no more than two electrons, a feature sometimes pictured as if the electron were spinning and where only two spin orientations were allowed: clockwise and counterclockwise.

The predicted sequence of electron filling might be best illustrated by looking at some examples. Hydrogen is 1s 1, the superscript referring to the number of electrons in the 1s substate. Lithium (three electrons) and sodium (eleven electrons) are 1s22s1 and 1s22s22p63s1, respectively. The latter configuration, for example, corresponds to one pair of 1s- electrons, one pair of

2s- electrons, three pairs of 2p- electrons (six total), and a final 3s -electron. Neon and argon are 1s22s22p6 and 1s22s22p63s23p6, respectively. They complete the horizontal periods of length eight.

The periods in which the d substates are filling are known as the d- block elements or transition metals . These ten elements increase the period length to eighteen elements. Some new Periodic Tables have adopted the convention of numbering the columns one through eighteen as a result.

The f- block, whose existence was recognized by American chemist Glenn Seaborg, has two rows containing nearly one-quarter of all the elements. The first row is known as the rare earth elements or lanthanides . The second f- block row is referred to as the actinides. The most common form of the Periodic Table, the Mendeleev-Seaborg form, has the f- elements at the bottom. Fourteen f- block elements increase the period length to thirty-two.

For nearly three centuries, a new element has been discovered every two and-one-half years, on average. Undoubtedly, more will be found. Although their names and their discoveries will likely involve controversies, their place at the table is already set.

see also Alchemy; Avogadro, Amedeo; Becquerel, Antoine-Henri; Bohr, Niels; Cannizzaro, Stanislao; Dalton, John; Lavoisier, Antoine; Mendeleev, Dimitri; Meyer, Lothar; Pauli, Wolfgang; Ramsay, William; RÖntgen, Wilhelm; Rutherford, Ernest; Seaborg, Glenn Theodore; Thomson, Joseph John.

Paul J. Karol

Bibliography

Marshall, James L. (2000). "A Living Periodic Table." Journal of Chemical Education 77:979983.

Mazurs, Edward G. (1974). Graphic Representations of the Periodic System during One Hundred Years, revised 2nd edition. University: University of Alabama Press.

van Spronsen, J. W. (1969). The Periodic System of Chemical Elements: A History of the First Hundred Years. New York: Elsevier.

Internet Resources

Winter, Mark. "WebElements Period Table." The University of Sheffield and WebElements Ltd., U.K. Available from <http://www.webelements.com>.

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Periodic Table

Periodic table

The periodic table is a chart that shows the chemical elements and their relationship to each other. The periodic table is a graphic way of representing the periodic law.

History of the periodic law

By the middle of the nineteenth century, about 50 chemical elements were known. One of the questions chemists were asking about those elements was the following: Is every element entirely different from every other element? Or are some elements related to other elements in some way? Are there patterns among the elements?

A number of chemists suggested various patterns. German chemist Johann Wolfgang Döbereiner (17801849) observed in 1829, for example, that three of the so-called halogen elements (chlorine, bromine, and iodine) could be classified according to their atomic weights. The atomic weight of bromine (79.9) turned out to be almost the exact average of the atomic weights of chlorine (35.5) and iodine (127), with 35.5 + 127 ÷ 2 = 81.25 (almost 79.9)

Most of these classification schemes were not very successful. Then, in about 1869, two chemists made almost the same discovery at almost the same time. Russian chemist Dmitry Mendeleev (18341907) and German chemist Julius Lothar Meyer (18301895) suggested arranging the elements according to their atomic weights. In doing so, Mendeleev and Meyer pointed out, the properties of the elements appear to recur in a regular pattern.

Words to Know

Atomic number: The number of protons in the nucleus of an atom; the number that appears over the element symbol in the periodic table.

Atomic weight: The average weight of all isotopes of a given element, expressed in units known as atomic mass units (amu).

Element: A pure substance that cannot be changed chemically into a simpler substance.

Family: A group of elements in the same column of the periodic table or in closely related columns of the table. (See Group.)

Group: A vertical column of the periodic table that contains elements possessing similar chemical characteristics. (See Family.)

Isotopes: Two or more forms of the same element with the same number of protons but different numbers of neutrons in the atomic nucleus.

Nucleus: The small core at the center of an atom that contains protons and (usually) neutrons.

Period: A horizontal row of elements in the periodic table. (See Row.)

Row: A horizontal set of elements in the periodic table. (See Period.)

Today, Mendeleev is usually given credit for discovering the periodic law because he took one step that Meyer did not. When all the elements are laid out in a table, some gaps appear. The reason for those gaps, Mendeleev said, was that other elements belonged there. But those elements had not yet been discovered.

Mendeleev went even further. He predicted the properties of those yet-to-be-discovered elements. He knew where they belonged in the periodic table, so he knew what elements they would be like. Remarkably, three of the elements Mendeleev predicted were discovered less than a decade after the periodic law was announced.

The modern periodic table

The periodic table used today is shown in Figure 1. It contains all of the known elements from the lightest (hydrogen: H) to the heaviest (meitnerium: Mt). Currently, there are 114 known elements, ranging from hydrogen, whose atoms have only one electron, to the as-yet unnamed element whose atoms contain 114 electrons. Each element has its own box in the periodic table. As shown in the sample at the top of the table, that box usually contains four pieces of information: the element's name, its symbol, its atomic number, and its atomic weight.

The table is divided in two directions, by rows and by columns. There are seven rows, called periods, and 18 columns, called groups or families. Two different numbering systems are used for the groups, as shown at the top of the table. The system using Roman numerals (IA, IIA, IIB, IVB, etc.) has traditionally been popular in the United States. The other system (1, 2, 3, 4, etc.) has traditionally been used in Europe and, a few years ago, was recommended for use in the United States as well.

Chemical elements in the same group tend to have similar chemical properties. Those in the same row have properties that change slowly from one end of the row to the other end. Figure 2 shows how one propertyatomic radiuschanges for certain elements in the table.

The appearance of the periodic table in Figure 1 is a little bit misleading, as is the case in almost every periodic table that is published. The reason for this misrepresentation is that two groups of elements shown at the bottom of the table actually belong within it. The Lanthanides, for example, belong in row 6 between lanthanum (#57) and hafnium (#72). Also, the Actinides belong in row 7 between actinium (#89) and unnilquadium (#104). The reason you don't see them there is that they simply don't fit. If they were actually inserted where they belong, the table would be much too wide to fit on a piece of paper or a wall chart. Thus, they are listed at the bottom of the table.

The diagonal line at the right of the table separates the elements into two major groups, the metals and nonmetals. Elements to the left of this line tend to be metals, while those to the right tend to be nonmetals. The elements that lie directly on the diagonal line are metalloidselements that behave sometimes like metals and sometimes like nonmetals.

The periodic table is one of the most powerful tools available to chemists and to chemistry students. Simply by knowing where an element is on the table, one can know a great deal about its physical and chemical properties.

Recent and future research

Recently, several man-made "superheavy" elements have been discovered. These include elements 110 and 111, both of which were made in late 1994 by an international team of scientists. 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. Elements 116 and 118 were recently discovered at a Berkeley, California, laboratory.

Other superheavy elements that have been predicted to exist have yet to be made in the laboratory, although research continues into the creation of these elements. Many exciting discoveries remain to be uncovered concerning the creation of new elements. With the periodic table as a guide, their place is already waiting for them.

[See also Atom; Atomic mass; Element, chemical ]

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periodic table

periodic table, chart of the elements arranged according to the periodic law discovered by Dmitri I. Mendeleev and revised by Henry G. J. Moseley. In the periodic table the elements are arranged in columns and rows according to increasing atomic number (see the table entitled Periodic Table).

There are 18 vertical columns, or groups, in the standard periodic table. At present, there are three versions of the periodic table, each with its own unique column headings, in wide use. The three formats are the old International Union of Pure and Applied Chemistry (IUPAC) table, the Chemical Abstract Service (CAS) table, and the new IUPAC table. The old IUPAC system labeled columns with Roman numerals followed by either the letter A or B. Columns 1 through 7 were numbered IA through VIIA, columns 8 through 10 were labeled VIIIA, columns 11 through 17 were numbered IB through VIIB and column 18 was numbered VIII. The CAS system also used Roman numerals followed by an A or B. This method, however, labeled columns 1 and 2 as IA and IIA, columns 3 through 7 as IIIB through VIB, column 8 through 10 as VIII, columns 11 and 12 as IB and IIB and columns 13 through 18 as IIIA through VIIIA. However, in the old IUPAC system the letters A and B were designated to the left and right part of the table, while in the CAS system the letters A and B were designated to the main group elements and transition elements respectively. (The preparer of the table arbitrarily could use either an upper-or lower-case letter A or B, adding to the confusion.) Further, the old IUPAC system was more frequently used in Europe while the CAS system was most common in America. In the new IUPAC system, columns are numbered with Arabic numerals from 1 to 18. These group numbers correspond to the number of s,p, and d orbital electrons added since the last noble gas element (in column 18). This is in keeping with current interpretations of the periodic law which holds that the elements in a group have similar configurations of the outermost electron shells of their atoms. Since most chemical properties result from outer electron interactions, this tends to explain why elements in the same group exhibit similar physical and chemical properties. Unfortunately, the system fails for the elements in the first 3 periods (or rows; see below). For example, aluminum, in the column numbered 13, has only 3 s,p, and d orbital electrons. Nevertheless, the American Chemical Society has adopted the new IUPAC system.

The horizontal rows of the table are called periods. The elements of a period are characterized by the fact that they have the same number of electron shells; the number of electrons in these shells, which equals the element's atomic number, increases from left to right within each period. In each period the lighter metals appear on the left, the heavier metals in the center, and the nonmetals on the right. Elements on the borderline between metals and nonmetals are called metalloids.

Group 1 (with one valence electron) and Group 2 (with two valence electrons) are called the alkali metals and the alkaline-earth metals, respectively. Two series of elements branch off from Group 3, which contains the transition elements, or transition metals; elements 57 to 71 are called the lanthanide series, or rare earths, and elements 89 to 103 are called the actinide series, or radioactive rare earths; a third set, the superactinide series (elements 122–153), is predicted to fall outside the main body of the table, but none of these has yet been synthesized or isolated. The nonmetals in Group 17 (with seven valence electrons) are called the halogens. The elements grouped in the final column (Group 18) have no valence electrons and are called the inert gases, or noble gases, because they react chemically only with extreme difficulty.

In a relatively simple type of periodic table, each position gives the name and chemical symbol for the element assigned to that position; its atomic number; its atomic weight (the weighted average of the masses of its stable isotopes, based on a scale in which carbon-12 has a mass of 12); and its electron configuration, i.e., the distribution of its electrons by shells. The only exceptions are the positions of elements 103 through 118; complete information on these elements has not been compiled. Larger and more complicated periodic tables may also include the following information for each element: atomic diameter or radius; common valence numbers or oxidation states; melting point; boiling point; density; specific heat; Young's modulus; the quantum states of its valence electrons; type of crystal form; stable and radioactive isotopes; and type of magnetism exhibited by the element (paramagnetism or diamagnetism).

See P. W. Atkins, The Periodic Kingdom: A Journey into the Land of Chemical Elements (1997).

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periodic table

periodic table Arrangement of the chemical elements in order of their atomic number in accordance with the periodic law first stated by Russian chemist Dmitri Mendeleyev in 1869. The metallic transition elements are arranged in the middle of the table between groups II and III. Alkali metals are in group I, and alkaline-earth metals in Group II. Metalloids and nonmetals are found from groups III to VII, with the halogens in group VII and the noble gases (inert gases) collected into group 0. The elements in each group have the same number of valence electrons and accordingly have similar chemical properties. Elements in the same horizontal period have the same number of electron shells.

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periodic table

pe·ri·od·ic ta·ble / ˌpi(ə)rēˈädik/ • n. Chem. a table of the chemical elements arranged in order of atomic number, usually in rows, so that elements with similar atomic structure (and hence similar chemical properties) appear in vertical columns.

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periodic table

periodic table a table of the chemical elements arranged in order of atomic number, usually in rows, so that elements with similar atomic structure (and hence similar chemical properties) appear in vertical columns.

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Periodic Table

Periodic Table

Construction of the table

Mendeleevs predictions

Layout of the periodic table

Electronic structure

Other trends

Recent and future research

Names of the elements

Resources

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 Cannizzaros 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.

Mendeleevs 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 Mendeleevs 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 tables 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.

Moseleys 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 Moseleys 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 Mendeleevs time the character of the periodic table has remained basically unchanged, providing testament to the power of his original insight.

After Moseleys 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 atoms 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 elements 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 (19071991) 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 (19121999), 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 atoms 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 (18791955) and Swedish inventor Alfred Nobel (18331896).

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 elements 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. Natures 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

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Periodic Table

Periodic table

The arrangement of the chemical elements into periods (horizontal rows) and groups (vertical columns) is called the periodic table. 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 the 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 130 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 until a time past the launching of the space shuttle .


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 Hennig Brand, a German scientist who discovered the element phosphorous in 1649. There were 63 known elements in 1869, the year Mendeleev created the periodic table.

Dmitri Ivanovitch Mendeleev (1834-1907) 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.

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). Mendeleev predicted the atomic mass of each element along with compounds they each should form. Within 15 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, were problems 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.

Refinement in the measurements of atomic mass, the ordering of the elements based on atomic number rather than atomic mass by Henry G. Moseley (1887-1915) in 1913, and the discovery of new elements have led to the continuing evolution of the periodic table. But since Mendeleev's time the periodic table has remained basically unchanged, providing testament to the power of his original insight.


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 us all through our everyday lives. From experience we 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 nonmetal elements familiar to us 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 nonmetals 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 we encounter the metalloids (or semimetals), 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 which 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-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 = 1.66X10-27 kilogram), 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. We will not go into detail on the physical meaning of subshells, we only need to know that 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 subshells, 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 sub-shell, 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 let us look at the first three elements of Group 16, oxygen, sulfur, and selenium. 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 we expect them to have similar chemistry. We are all familiar with the compound water which has the formula H2O. Likewise there are the compounds H2S and H2Se. In a similar fashion, if you are told that the Group 15 element nitrogen, with the valence shell electronic configuration s2p3, forms the compound NH3 (called ammonia ), can you infer the formula of the compound that forms between phosphorus and hydrogen? By analogy to NH3 we 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 which 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) we find a positively charged sodium atom (Na+) and a negatively charged chlorine atom (Cl-). If we look at the valence shell electron configuration of each ion we 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 manmade, 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 manmade 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 E. M. Macmillan 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. Glenn T. Seaborg, 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 104-111, 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 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 manmade "superheavy" 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 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. Elements 116 and 118 were recently discovered at a Berkeley laboratory. Other superheavy elements which have been predicted to exist have yet to be made 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, stannum, 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 einsteinium (Es) and nobelium (No), named after Albert Einstein and Alfred Nobel.

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 104-109 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 which appear on the current periodic table are, in order, rutherfordium (Rf), dubnium (Db), seaborgium, bohrium (Bh), hassium (Hs), and meitnerium (Mt).

A particular controversy among these elements involved element 106 which 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 H. The Norton History of Chemistry. New York: W.W. Norton & Company Inc., 1992.

Hoffmann, Roald, and Torrence, Vivian. Chemistry ImaginedReflections on Science. Washington: Smithsonian Institutional Press, 1993.

Roberts, Royston M. Serendipity: Accidental Discoveries inScience. New York: John Wiley & Sons Inc., 1989.


periodicals

Naeye, Robert. "An Island of Stability." Discover (August 1994): 15.


Michael G. Roepel

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.

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