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 shell—a 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 (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 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 (1834–1907) returned from the meeting to St. Petersburg, Russia, he pondered Cannizzaro's list of atomic weights along with an immense amount of information he had gathered about the properties of elements. Mendeleyev found that when he arranged the elements in order of increasing atomic weight, similar properties were repeated at regular intervals—they displayed periodicity. 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 (1856–1940) in England led to the discovery of the electron and, later, the proton. In 1932, James Chadwick (1891–1974), 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 (1887–1915), 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 (1875–1946).
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 elements—lithium, sodium, potassium, rubidium, cesium and francium—all 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 universe—only 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 table—groups 1, 2, 17, and 18, the alkali metals, the alkaline earths, the halogens and the noble gases—enables 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 not—they 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 3–12, including the lanthanide and actinide elements, are called the transition elements or transition metals. The other elements, Groups 1,2, and 13–18, 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
"Periodic Table (Predicting the Structure and Properties of the Elements)." World of Earth Science. . Encyclopedia.com. 27 Jun. 2017 <http://www.encyclopedia.com>.
"Periodic Table (Predicting the Structure and Properties of the Elements)." World of Earth Science. . Encyclopedia.com. (June 27, 2017). http://www.encyclopedia.com/science/encyclopedias-almanacs-transcripts-and-maps/periodic-table-predicting-structure-and-properties-elements
"Periodic Table (Predicting the Structure and Properties of the Elements)." World of Earth Science. . Retrieved June 27, 2017 from Encyclopedia.com: http://www.encyclopedia.com/science/encyclopedias-almanacs-transcripts-and-maps/periodic-table-predicting-structure-and-properties-elements
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 substances—earth, air, fire, and water—of 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.
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.
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 (1837–1898)
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.
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.
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.
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 found—called an orbital because it replaced Bohr's planetary fixed orbit idea—could 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
Marshall, James L. (2000). "A Living Periodic Table." Journal of Chemical Education 77:979–983.
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.
Winter, Mark. "WebElements Period Table." The University of Sheffield and WebElements Ltd., U.K. Available from <http://www.webelements.com>.
"Periodic Table." Chemistry: Foundations and Applications. . Encyclopedia.com. (June 27, 2017). http://www.encyclopedia.com/science/news-wires-white-papers-and-books/periodic-table
"Periodic Table." Chemistry: Foundations and Applications. . Retrieved June 27, 2017 from Encyclopedia.com: http://www.encyclopedia.com/science/news-wires-white-papers-and-books/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 (1780–1849) 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 (1834–1907) and German chemist Julius Lothar Meyer (1830–1895) 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 property—atomic radius—changes 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 metalloids—elements 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 ]
"Periodic Table." UXL Encyclopedia of Science. . Encyclopedia.com. (June 27, 2017). http://www.encyclopedia.com/science/encyclopedias-almanacs-transcripts-and-maps/periodic-table
"Periodic Table." UXL Encyclopedia of Science. . Retrieved June 27, 2017 from Encyclopedia.com: http://www.encyclopedia.com/science/encyclopedias-almanacs-transcripts-and-maps/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).
"periodic table." The Columbia Encyclopedia, 6th ed.. . Encyclopedia.com. (June 27, 2017). http://www.encyclopedia.com/reference/encyclopedias-almanacs-transcripts-and-maps/periodic-table
"periodic table." The Columbia Encyclopedia, 6th ed.. . Retrieved June 27, 2017 from Encyclopedia.com: http://www.encyclopedia.com/reference/encyclopedias-almanacs-transcripts-and-maps/periodic-table
"periodic table." World Encyclopedia. . Encyclopedia.com. (June 27, 2017). http://www.encyclopedia.com/environment/encyclopedias-almanacs-transcripts-and-maps/periodic-table
"periodic table." World Encyclopedia. . Retrieved June 27, 2017 from Encyclopedia.com: http://www.encyclopedia.com/environment/encyclopedias-almanacs-transcripts-and-maps/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.
"periodic table." The Oxford Pocket Dictionary of Current English. . Encyclopedia.com. (June 27, 2017). http://www.encyclopedia.com/humanities/dictionaries-thesauruses-pictures-and-press-releases/periodic-table-0
"periodic table." The Oxford Pocket Dictionary of Current English. . Retrieved June 27, 2017 from Encyclopedia.com: http://www.encyclopedia.com/humanities/dictionaries-thesauruses-pictures-and-press-releases/periodic-table-0
"periodic table." The Oxford Dictionary of Phrase and Fable. . Encyclopedia.com. (June 27, 2017). http://www.encyclopedia.com/humanities/dictionaries-thesauruses-pictures-and-press-releases/periodic-table
"periodic table." The Oxford Dictionary of Phrase and Fable. . Retrieved June 27, 2017 from Encyclopedia.com: http://www.encyclopedia.com/humanities/dictionaries-thesauruses-pictures-and-press-releases/periodic-table