The smallest unit of a chemical element that has the properties of that element is called an atom. Many elements (e.g., helium) occur as single atoms. Other elements occur as molecules made up of more than one atom. Elements that ordinarily occur as diatomic molecules include hydrogen, nitrogen, oxygen, and the halogens, but oxygen also occurs as a triatomic form called ozone. Phosphorus usually occurs as a tetratomic molecule, and crystalline sulfur occurs as molecules containing eight atoms.Atomic Number and Mass Number
Regardless of how many atoms the element is composed of, each atom has the same number of protons in its nucleus, and this is different from the number in the nucleus of any other element. Thus this number, called the atomic number (at. no.), defines the element. For example, the element carbon consists of atoms all with at. no. 6, i.e., all having 6 protons in the nucleus; any atom with at. no. 6 is a carbon atom. By 2006, 117 elements were known, ranging from hydrogen with an at. no. of 1 to an as yet unnamed element (temporarily known as ununoctium) with an at. no. of 118. (See the table entitled Elements for an alphabetical list of all the elements, including their symbols, atomic numbers, atomic weights, and melting and boiling points.) The nuclei of most atoms also contain neutrons. The total number of protons and neutrons in the nucleus of an atom is called the mass number. For example, the mass number of a carbon atom with 6 protons and 6 neutrons in its nucleus is 12.Isotopes
Although all atoms of an element have the same number of protons in their nuclei, they may not all have the same number of neutrons. Atoms of an element with the same mass number make up an isotope of the element. All known elements have isotopes; some have more than others. Hydrogen, for example, has only 3 isotopes, while xenon has 16. Approximately 300 naturally occurring isotopes are known, and more than 2,500 radioactive isotopes have been artificially produced (see synthetic elements). There are 13 isotopes of carbon, having from 2 to 14 neutrons in the nucleus and therefore mass numbers from 8 to 20.
Not all of the elements have stable isotopes. Some have only radioactive isotopes, which decay to form other isotopes, usually of other elements (see radioactivity). In some cases all the isotopes of an element are very unstable, and the element is therefore not found in nature. Only 94 of the elements are known to occur naturally on earth. Of these, 6 occur in minute amounts produced by the decay of other elements. These 6 extremely scarce elements and those that do not occur at all naturally were discovered when they were produced in the laboratory; they are often called the man-made, artificially produced, or synthetic elements.Atomic Mass and Atomic Weight
Atoms are not very massive; a carbon atom weighs about 2 × 10-23 grams. Because atoms have so little mass, a unit much smaller than the gram is used. In the current system (adopted in 1960-61) the unit of atomic mass, called atomic mass unit (amu), is defined as exactly 1/12 the mass of an atom of carbon-12. The atomic weight of an element is the mean (weighted average) of the atomic masses of all the naturally occurring isotopes. Carbon has two principal naturally occurring isotopes, carbon-12 and carbon-13. Carbon-12, whose mass is defined as exactly 12 amu, constitutes 98.89% of naturally occurring carbon; carbon-13, whose mass is 13.00335 amu, constitutes 1.11%. (There are also small traces of the radioactive isotope carbon-14.) The atomic weight of the element is determined by multiplying the percent abundance of each isotope by the atomic mass of the isotope, adding these products, and dividing by 100. However, isotope abundance is often determined by the medium of the source, solid, liquid, or gas, and the average atomic weight may fluctuate. Thus, for carbon, [(98.89 × 12.000) + (1.11 × 13.00335)]/100 = 12.01115, which is the atomic weight of the element carbon in amu. Certain synthetic elements exist only momentarily in the form of a few short-lived isotopes; in such cases the concept of atomic weight cannot be applied.
Properties of an element are sometimes classed as either chemical or physical. Chemical properties are usually observed in the course of a chemical reaction, while physical properties are observed by examining a sample of the pure element. The chemical properties of an element are due to the distribution of electrons around the atom's nucleus, particularly the outer, or valence, electrons; it is these electrons that are involved in chemical reactions. A chemical reaction does not affect the atomic nucleus; the atomic number therefore remains unchanged in a chemical reaction.
Some properties of an element can be observed only in a collection of atoms or molecules of the element. These properties include color, density, melting point, boiling point, and thermal and electrical conductivity. While some of these properties are due chiefly to the electronic structure of the element, others are more closely related to properties of the nucleus, e.g., mass number.
The elements are sometimes grouped according to their properties. One major classification of the elements is as metals, nonmetals, and metalloids. Elements with very similar chemical properties are often referred to as families; some families of elements include the halogens, the inert gases, and the alkali metals. In the periodic table the elements are arranged in order of increasing atomic weight in such a way that the elements in any column have similar properties.
Each element is assigned an official symbol by the International Union of Pure and Applied Chemistry (IUPAC). For example, the symbol for carbon is C, and the symbol for silver is Ag [Lat. argentum = silver]. There are several ways of designating an isotope. One designation consists of the name or symbol of the element followed by a hyphen and the mass number of the isotope; thus the isotope of carbon with mass number 12 can be designated carbon-12 or C-12. The mass number is often written as a superscript, e.g., C12; sometimes the atomic number is written as a subscript preceding the symbol, e.g., 6C12. The IUPAC rules for nomenclature of inorganic chemistry state that the subscript atomic number and superscript mass number should both precede the symbol, e.g., 126C.
Many isotopes were given special names and symbols when they were first discovered in natural radioactive decay series (e.g., uranium-235 was called actinouranium and represented by the symbol AcU). This practice is discouraged in the modern nomenclature except in the case of hydrogen. The isotopes hydrogen-2 and hydrogen-3 are usually called deuterium and tritium, respectively. Hydrogen-1, the most abundant isotope, has the name protium but is usually simply called hydrogen. Newly discovered elements that have been synthesized by one laboratory and not yet confirmed by a second are given a provisional name based on Greek and Latin roots; when the discovery is confirmed, the laboratory that first made it may suggest a name for the element.
Some elements have been known since antiquity. Gold ornaments from the Neolithic period have been discovered. Gold, iron, copper, lead, silver, and tin were used in Egypt and Mesopotamia before 3000 B.C. However, recognition of these metals as chemical elements did not occur until modern times.Greek Concept of the Elements
The Greek philosophers proposed that there are basic substances from which all things are made. Empedocles proposed four basic "roots," earth, air, fire, and water, and two forces, harmony and discord, joining and separating them. Plato called the roots stoicheia (elements). He thought that they assume geometric forms and are made up of some more basic but undefined matter. A different theory, that of Leucippus and his followers, held that all matter is made up of tiny indivisible particles (atomos).
This theory was rejected by Aristotle, who expanded on Plato's theory. Aristotle believed that different forms (eidos) were assumed by a basic material, which he called hulé. The hulé had four basic properties, hotness, coldness, dryness, and moistness. The four elements differ in their embodiment of these properties; fire is hot and dry, earth cold and dry, water cold and moist, and air hot and moist. Although Aristotle proposed that an element is "one of those simple bodies into which other bodies can be decomposed and which itself is not capable of being divided into others," he thought the metals to be made of water, and called mercury "silver water" (chutos arguros). His idea that matter was a single basic substance that assumed different forms led to attempts by the alchemists to transmute other metals into gold.Evolution of Modern Concepts
Although much early work was done in chemistry, especially with metals, and many recipes were recorded, there were few developments in the conception of the elements. In the 16th cent. Paracelsus proposed salt, mercury, and sulfur as three "principles" of which bodies were made, although he apparently also believed in the four "elements." Van Helmont (c.1600) rejected the four elements and three principles, substituting two elements, air and water.
Robert Boyle rejected these early theories and proposed a definition of chemical elements that led to the currently accepted definition. His definition is strikingly similar to Aristotle's earlier definition. In The Sceptical Chymist (1661) Boyle wrote, "I now mean by elements … certain primitive and simple, or perfectly unmingled bodies; which not being made of any other bodies, or of one another, are the ingredients of which all those called perfectly mixed bodies [chemical compounds] are immediately compounded, and into which they are ultimately resolved."
Whereas Aristotle and other early philosophers tried to determine the identity of the elements solely by reason, Boyle and later scientists used the results of numerous experiments to identify the elements. In 1789 Antoine Lavoisier published a list of chemical elements based on Boyle's definition; this encouraged adoption of standard names for the elements. Although some of his elements are now known to be compounds, such as metallic oxides and salts, they were at the time accepted as elements since they could not be decomposed by any method then known.
In 1803 John Dalton proposed (as part of his atomic theory) that all atoms of an element have identical properties (including mass), that these atoms are unchanged by chemical action, and that atoms of different elements react with one another in simple proportions. Although symbols for some of the elements already existed, they were by no means universally accepted, and each compound also had a unique symbol that was unrelated to its chemical composition. Dalton devised a new set of circular symbols for the elements and used a combination of elemental symbols to represent a compound. For example, his symbol for oxygen was ○, and for hydrogen ⊙. Since he thought water contained one atom of hydrogen for every atom of oxygen, he formed the symbol for water by writing the symbols for hydrogen and oxygen touching one another, ⊙&nosp;○. J. J. Berzelius was the first to use the modern method, letting one or two letters of the element's name serve as its symbol. He also published an early table of atomic weights of 24 elements with most values very close to those now in use.Discovery of the Elements
As noted above, some of the elements were discovered in prehistoric times but were not recognized as elements. Arsenic was discovered around 1250 by Albertus Magnus, and phosphorus was discovered about 1674 by Hennig Brand, an alchemist, who prepared it by distilling human urine. Only 12 elements were known before 1700, and only about twice that many by 1800, but by 1900 over 80 elements had been identified. In 1919 Ernest Rutherford found that hydrogen was given off when nitrogen was bombarded with alpha particles. This first transmutation encouraged further study of nuclear reactions, and eventually led to the discovery in 1937 of technetium, the first synthetic element. Neptunium (atomic number 93) was the first transuranium element to be synthesized (1940). Its discovery prompted the search that led to the discovery of other transuranium elements.
See J. Emsley, The Elements (1991); A. Swertka, A Guide to the Elements (1996); P. W. Atkins, The Periodic Kingdom (1997); N. N. Greenwood and A. Earnshaw, Chemistry of the Elements (2d ed. 1997).
Any of the chemical elements after uranium in the periodic table (with atomic numbers greater than 92). All are radioactive (see radioactivity), with half-lives ranging from tens of millions of years to fractions of a millisecond. Only two, neptunium (93) and plutonium (94), occur in nature, and only as traces in uranium ores as a result of neutron irradiation. Transuranium elements with atomic numbers through 116, along with 118, have been produced in laboratories. Each appears to resemble the elements above it in the periodic table; in particular, the actinides, thorium (90) through lawrencium (103), are similar to the lanthanides, cerium (58) through lutetium (71). The naming of the transuranium elements has been fraught with controversy regarding which laboratory first made the discovery and should propose the name and whether elements should be named for living persons. Seealso Glenn Seaborg.
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Any chemical element with valence electrons in two shells instead of only one. This structure gives them their outstanding ability to form ions containing more than one atom (complex ions, or coordination compounds), with a central atom or ion (often of a transition metal) surrounded by ligands in a regular arrangement. Theories on the bonding in these ions are still being refined. The elements in the periodic table from scandium to copper (atomic numbers 21–29), yttrium to silver (39–47), and lanthanum to gold (57–79, including the lanthanide series) are frequently designated the three main transition series. (Those in the actinide series and beyond, 89–111, also qualify.) All are metals, many of major economic or industrial importance (e.g., iron, gold, nickel, titanium). Most are dense, hard, and brittle, conduct heat and electricity well, have high melting points, and form alloys with each other and other metals. Their electronic structure lets them form compounds at various valences. Many of these compounds are coloured and paramagnetic (see paramagnetism) and (as do the metals themselves) often act as catalysts. Seealso rare earth metal.
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Any of a large class of chemical elements including scandium (atomic number 21), yttrium (39), and the 15 elements from 57 (lanthanum) to 71 (see lanthanides). The rare earths themselves are pure or mixed oxides of these metals, originally thought to be quite scarce; however, cerium, the most plentiful, is three times as abundant as lead in the Earth's crust. The metals never occur free, and the pure oxides never occur in minerals. These metals are similar chemically because their atomic structures are generally similar; all form compounds in which they have valence 3, including stable oxides, carbides, and borides.
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Any of the 19 chemical elements that occur as minerals and are found in nature uncombined with other elements. They are commonly divided into three groups: metals (platinum, iridium, osmium, iron, zinc, tin, gold, silver, copper, mercury, lead, chromium); semimetals (bismuth, antimony, arsenic, tellurium, selenium); and nonmetals (sulfur, carbon). Members of the group of native elements form under widely varying physicochemical conditions and in very different rock types. Many deposits are sufficiently abundant to be commercially important.
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Any of five nonmetallic elements—fluorine, chlorine, bromine, iodine, and astatine—with similar chemical properties. They occur in the second rightmost column of the periodic table as usually arranged. All are highly reactive oxidizing agents (see oxidation-reduction) with valence 1 (for fluorine, the only valence). They combine readily with most metals and nonmetals to form a variety of compounds and never occur uncombined in nature. A radioactive element, astatine occurs naturally in minute amounts as an intermediate decay product; it has no stable nonradioactive isotopes. Halogen salts formed with metal atoms (halides) are very stable; sodium chloride is the most familiar. The halogen lamp takes its name from the halogens included in the gas within its tungsten-filament bulb, added to prolong filament life and increase brightness.
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One of the 117 presently known kinds of substances that constitute all matter at and above the level of atoms (the smallest units of any element). All atoms of an element are identical in nuclear charge (number of protons) and number of electrons (see atomic number), but their mass (atomic weight) may differ if they have different numbers of neutrons (see isotope). Each permanently named element has a one- or two-letter chemical symbol. Elements combine to form a wide variety of compounds. All elements with atomic numbers greater than 83 (bismuth), and some isotopes of lighter elements, are unstable and radioactive (see radioactivity). The transuranium elements, with atomic numbers greater than 92 (see uranium), artificially created by bombardment of other elements with neutrons or other particles, were discovered beginning in 1940. The most common elements (by weight) in Earth's crust are oxygen, 49percnt; silicon, 26percnt; aluminum, 8percnt; and iron, 5percnt. Of the known elements, 11 (hydrogen, nitrogen, oxygen, fluorine, chlorine, and the six noble gases) are gases under ordinary conditions, two (bromine and mercury) are liquids (two more, cesium and gallium, melt at about or just above room temperature), and the rest are solids. Seealso periodic table.
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Any of the series of 15 consecutive chemical elements in the periodic table from actinium to lawrencium (atomic numbers 89–103). All are radioactive heavy metals; and only the first four (actinium, thorium, protactinium, and uranium) occur in nature in appreciable quantities. The other 11 (the transuranium elements) are unstable and are produced only artificially. Actinides are transition elements, so their atoms have similar configurations and similar physical and chemical behaviour; the most usual valences are 3 and 4.
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