A chemical element is a sample of matter that cannot be separated into simpler species by chemical or physical means (excluding nuclear reactions). ELEMENTS AND COMPOUNDS Most matter is found in nature as mixtures, which chemists can separate into pure substances (compounds and elements). A chemical compound is a sample of matter that contains two or more elements in the same proportions throughout. Every chemical compound can be described by a chemical formula. Compounds can be broken down chemically into their constituent elements. (Note that the term substance refers only to elements and compounds, that is, to samples of matter in which the proportion of each element is fixed and reproducible. The term element refers to an elemental substance, that is, one that cannot be separated into simpler substances. The term compound refers to a compound substance.) ATOMIC STRUCTURE Atomic structure distinguishes one element from another. An atom is composed of a nucleus containing positively charged particles called protons (and particles with no charge, called neutrons), surrounded by negatively charged particles called electrons. The number of electrons equals the number of protons, so that the atom as a whole has no charge. Each element is characterized by the number of protons that an atom of that element contains, called its atomic number. All atoms having the same atomic number have the same chemical properties; for example, all atoms with six protons have the properties of the element called carbon. Besides its name, each element is given a symbol; the symbol for carbon is C. ELEMENTS OF THE ANCIENTS The concept of an element as it is understood today did not develop until the 17th century. In the days of the classical Greek philosophers \TEmpedocles\t, \TPlato\t, and \TAristotle\t, the material world was considered to be composed of four or five fundamental substances, or "elements." Empedocles named earth, air, fire, and water as elemental substances, and Aristotle added ether (a perfect substance of which heavenly bodies are composed). These elements were not really material substances, however, but idealizations of material substances. Qualities of temperature (hot or cold) and moisture (wet or dry) were attributed to each of the four classical elements. Fire, for example, had the two adjacent qualities hot and dry. Several actual elements were known to ancient civilizations. Gold, and to a lesser extent silver and copper, are found in their uncombined, or elemental, form in nature and have been used as pure metals since 3000 BC and probably earlier. The ancients also made iron, tin, lead, mercury, copper, and silver by reducing their ores to their elemental forms by smelting, often with carbon in the form of charcoal. The chemical symbols of all of the metals known to the ancient world are derived from their Latin names: gold is Au from aurum, silver Ag from argentum, iron Fe from ferrum, tin Sn from stannum, lead Pb from plumbum, and mercury Hg from hydrargum. Sulfur and carbon, two nonmetals found uncombined in nature as solids, were also known to the ancients. Many large sulfur deposits are found at or near the Earth's surface. The ancients referred to sulfur (as well as other inflammable substances) as brimstone, and vast quantities of native sulfur are still mined for the production of inorganic chemicals such as sulfuric acid. Carbon is found in nature as diamond, graphite, and, in less pure form, charcoal. MEDIEVAL ALCHEMY Medieval alchemists prepared a number of elements from compounds or mixtures. They discovered and named arsenic, antimony, bismuth, and phosphorus, but they did not classify them--or the elements known to the ancients--as elements. They recognized, however, that certain material substances could not be changed into others. Alchemists could transmute many substances into others by means of what are now recognized as chemical reactions, but they never achieved their dream of transmuting base metals, such as copper and lead, into precious ones, such as silver and gold. ROBERT BOYLE The first challenge to the philosophical definition of elements came during the 17th century. Robert \TBoyle\t was a natural philosopher, an active empiricist who believed that experiments should precede and be the test of a theory. In The Sceptical Chymist (1661) he stated that only observable and weighable substances should be regarded as elements and that elements are the primary substances from which other substances are made. Although Boyle was a founder of \Tqualitative chemical analysis\t, a science that tries to identify the constituent substances of a mixture, he did not name a set of substances that he considered to be elements. He did believe, however, that elements are composed of fundamental particles, and he considered metals to be elements. SCIENTIFIC CLASSIFICATION AND DISCOVERY OF ELEMENTS Antoine \TLavoisier\t established a firm foundation for chemistry in the 18th century. Lavoisier recognized that weight (more fundamentally, mass) is a property of all matter and that matter (and therefore weight or mass) is conserved in chemical reactions. He also defined an element as a substance that is not decomposable into simpler substances by chemical analysis; therefore, elements must always increase their weight when they react to form compounds. In his Traite elementaire de chimie (Elementary Treatise of Chemistry, 1789) he listed 33 elements. Some, such as phosphorus, carbon, silver, and antimony, had been known since the days of the alchemists. Others, such as oxygen, hydrogen, and zinc, had been isolated during the 18th century. Lavoisier suspected that some of his elements were compounds that would later be decomposed into elements, and he was correct. For example, he listed chalk as an element; only a few decades later chalk (calcium oxide) was decomposed, and the constituent element, calcium, was isolated. Lavoisier also listed light and heat as elements, although he was troubled because they had no weight. Atomic Theory Lavoisier's ideas were confirmed experimentally and theoretically by several chemists, who showed that substances always react in fixed proportions by weight and derived tables of combining weights, and by John \TDalton\t, whose atomic theory extended the concept that matter is composed of fundamental particles called atoms. In 1810, Dalton stated that all atoms of a given element have the same mass. Dalton analyzed experimental data on combining weights and set forth the first table of atomic weights--a table of the elements with numbers indicating their weights relative to one another. Since all these weights were relative, a particular atomic weight had to be chosen as standard. Jons \TBerzelius\t, the most famous chemist of the early 19th century, isolated and identified several new elements and established an accurate table of atomic weights for 50 elements. (A few of Berzelius's atomic weights were in error by factors of 2 or 3 because he did not know the correct chemical formulas of their compounds.) Berzelius also proposed chemical symbols for elements by taking one or two letters from the name of each element. The Periodic Table By 1860 more than 60 elements were known. Several chemists recognized that elements could be grouped into families with similar chemical properties, and various schemes of classification of the elements were proposed. In 1869, Dmitry \TMendeleyev\t published a \Tperiodic table\t of the elements, grouping them in families and showing several vacancies between the elements known at that time. In 1871 he published a better table and boldly predicted the properties of three of the "missing" elements. Gallium, germanium, and scandium were discovered within 15 years of Mendeleyev's predictions, and their properties closely matched the properties predicted by Mendeleyev. Mendeleyev used atomic weights and chemical properties to arrange his periodic table of the elements. During the same periods, Lothar \TMeyer\t developed a nearly identical periodic table, which he used to show that many physical properties (atomic volumes, hardness, boiling points, and others) also varied in a regular way among the chemical families. Spectroscopy In 1854 an American physician, David \TAlter\t, used prisms to show that the spectra of vaporized alloys produced the characteristic spectral lines of each element in the alloy. Alter hypothesized that each element has a characteristic spectrum and that the spectra of astronomical bodies could be used to identify their constituent elements. Following the invention of the spectroscope (1859), it was demonstrated that the emission or absorption spectrum of a vaporized element is indeed a sensitive and specific physical property that could be used to identify it. Cesium was the first element to be discovered (1860) from its spectrum, when two previously unseen blue lines were found in the spectrum of a mineral water solution. (The name cesium is derived from the Latin caesius, the blue color of the sky.) Many other elements, including the inert gases, were thereafter identified from their characteristic atomic spectra. Helium, for example, was first observed as a new yellow line in a spectrum of the Sun taken during a solar eclipse in 1868; 27 years later an identical line was found in the spectrum of a gaseous residue obtained from a mineral containing uranium, confirming that helium is indeed an element, and present on the Earth. Radioactivity One year after helium was found in a uranium mineral, Antoine Henri \TBecquerel\t discovered that all uranium compounds emit radiations that darken photographic plates. The helium found in minerals is produced by the disintegration of uranium to thorium. The helium nuclei produced by this disintegration are called alpha particles; when neutralized by acquiring two electrons from neighboring atoms, they become helium gas, which remains trapped in the mineral. The search for radioactivity in minerals led to the discovery of a wide variety of radioactive elements, beginning with polonium and radium by Marie and Pierre \TCurie\t. By 1917 three additional radioactive elements (radon, actinium, and protactinium) had been discovered. In addition, many of the short-lived isotopes of the uranium and thorium decay series were incorrectly identified as new elements. Sensitive spectroscopic and radiochemical techniques led to many more incorrect claims for discovery of new elements. One spectroscopic technique, however, has proved invaluable in identifying elements. In 1913, Henry \TMoseley\t showed that each element emits characteristic X rays and, moreover, that there is a precise mathematical relationship between the X-ray wavelength of an element and its atomic number. Previously, when elements had been classified on the basis of their atomic weights, the heavier element in each of three pairs (argon and potassium, cobalt and nickel, and tellurium and iodine) preceded the lighter one in the periodic table. Moseley showed that the atomic number is the distinguishing property of an element. Almost all of the elements discovered in the 20th century are radioactive. Two, technetium and promethium, filled previous gaps in the periodic table and were discovered among the products of the nuclear fission of uranium. The others have atomic numbers higher than 92 (uranium). PHYSICAL PROPERTIES OF ELEMENTS Allotropes Many elements can exist in more than one form; the different forms are called \Lallotrope\ls. Sulfur has several allotropes; at room temperature, its most stable form is a crystal form called orthorhombic. Orthorhombic sulfur changes to monoclinic sulfur at 95.5 deg C; both of these allotropes contain 8-membered rings of sulfur atoms. Other allotropes of sulfur have 6-, 7-, 9-, 10-, or 12-membered rings or long chains of sulfur atoms. Other commonly seen allotropes include those of carbon (diamond and graphite) and tin (gray and white). Isotopes Early in the study of radioactivity, it was recognized that one radioactive decay product of radium had properties identical to lead. In 1914, Theodore W. \TRichards\t and Max E. Lembert announced that the atomic weight of lead from different radioactive ores varied by as much as 1%. In 1913, Frederick \TSoddy\t proposed that atoms of the same element, produced from radioactive disintegrations, might have different masses; he proposed the name \Lisotope\ls for two atoms of the same atomic number but different masses. Shortly after World War I, Francis W. \TAston\t developed the mass spectrograph, an instrument for determining atomic masses. Aston showed that many elements not produced by radioactivity also have isotopes. In 1930, James \TChadwick\t explained that almost all atomic nuclei contain neutrons, neutral particles that contribute mass but not charge to the nucleus. Isotopes are atoms of a particular element that have the same nuclear charge (number of protons) as the element but different numbers of neutrons. The number of protons plus neutrons in a given atom is its mass number. Modern mass spectrometers can measure atomic masses very precisely. Atomic Weights Atomic weights represent experimentally averaged atomic masses of all the naturally occurring isotopes of an element. The modern atomic weight scale is based on the convention that the mass of a C-12 atom is exactly 12 atomic mass units. Some elements, such as fluorine, have only one naturally occurring isotope; their atomic weights are measured with mass spectrometers to very high accuracy. The atomic weights of other elements are less accurately known because they are averages of all the naturally occurring isotopes of the given element. The isotopic composition (and, therefore, the atomic weight) of some elements varies according to the source or prior treatment of the sample. CHEMICAL PROPERTIES Metals and Nonmetals Most chemical elements are \Lmetal\ls: 75 of the first 103 elements are known to be metals in their elemental state. Francium (element 87) and elements 99 and beyond have chemical properties characteristic of metals, but they have such short HALF-LIVES that observable quantities have not been isolated in pure form. Metals are characterized by their weakly bound outermost electrons; these electrons' relative freedom of motion causes the characteristic conductivity, luster, and ductility of metals and results in the formation of positive ions (cations) by all metallic elements in ionic compounds. Seventeen elements are nonmetals. With the exception of the inert gases, they exist as diatomic or polyatomic molecules in their elemental form, share electrons in covalent bonds, and form negative ions (anions) in ionic compounds. Between the metals and the nonmetals are borderline elements called semimetals or metalloids. Electronegativity Electronegativity is the relative attraction of an atom for electrons in a covalent bond. Elements at the right and the top of the periodic table have the largest electronegativities; fluorine is the most electronegative element, and francium is the least electronegative. Elements with similar electronegativities form covalent compounds with electrons shared equally in bonds: nitrogen trichloride is a covalent compound. Elements with somewhat different electronegativities form polar covalent compounds, in which the electron cloud forming a bond is displaced slightly toward the more electronegative element: phosphorus trifluoride is a polar covalent compound in which the electron density is greater near the fluorine atoms. Elements with widely different electronegativities form ionic compounds, in which the electronegative atom forms an anion and the electropositive atom forms a cation. Salts, such as sodium chloride, are examples of ionic compounds. Many other properties, both chemical and physical, of an element can be classified and interpreted on the basis of the element's position in the periodic table. ORIGIN AND ABUNDANCE OF THE ELEMENTS At present the most plausible theory of the origin of the universe is that it formed from the explosion of a tiny, extremely dense fireball several billion years ago. During the first few seconds after this big bang (see \Tbig bang theory\t), the energy density was so great that only radiation and fundamental particles existed. As the universe cooled and expanded, after the big bang, however, hydrogen formed and after millions of years condensed into galaxies and then stars. The hydrogen in these first-generation stars burns by forming heavier nuclei and releasing energy, but even if such a star explodes in a \Tsupernova\t, the heaviest isotope produced is Fe-56. The gaseous products of the explosion of first-generation stars contain many light elements, and when the gases condense to form second-generation stars, many nuclear reactions can take place. Some of these nuclear reactions emit neutrons, which are captured by other nuclei to form heavier nuclei that decay by beta-emission to form elements with still higher atomic numbers. As a result, the universe contains small amounts of elements heavier than Fe-56, roughly in proportion to their nuclear stability; more than 60 elements have been identified in the Sun's spectrum. As the Sun formed, the disk-shaped rotating mass of the size of the solar system developed many turbulences. These turbulences led to the infrequent accretion of small masses of heavy elements, ultimately forming protoplanets surrounded by hydrogen-helium atmospheres. When the central mass, the Sun, began radiating energy as a result of its nuclear reactions, hydrogen and helium were driven off the inner planets. The Earth gradually evolved into an iron-nickel core, a silicate mantle, and a silicate crust and later developed an atmosphere and hydrosphere (the bodies of water and ice on the Earth's surface). In each of these regions, elemental abundances are different. For example, relative abundances of elements in the atmosphere and hydrosphere and on the Earth's surface are a combination of geological and biological processes. The abundances of elements in living things reflect their abundances in the Earth's crust and their usefulness in biochemical processes. LESTER R. MORSS Bibliography: Cox, P. A., The Elements (1989); Donahue, Jerry, Structure of the Elements (1974; repr. 1982); Emsley, J., The Elements (1989); Ruben, S., Handbook of the Elements (1985); Trifonov, D. N., Chemical Elements: How They Were Discovered (1985); Weeks, Mary, Discovery of the Elements, 7th ed. (1968).