Atom: Structure, Properties, and Impact on Modern Technology

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Introduction to Atom

Atom, the smallest part of a chemical element that retains the identity of the element. The study of the structure and behavior of the atom by chemists and physicists has contributed greatly to the scientific understanding of matter and its properties. The use of this knowledge has greatly affected present-day technology, leading to the commercial development of electronics and nuclear energy and to the production of many new kinds of materials and products, including plastics and synthetic fibers.

Atoms are so small that a hundred million of them, placed side by side, would make a row only about one inch (2.5 cm) long. Individual atoms can be viewed only with certain types of nonoptical microscopes that have extremely high magnifications.

Each chemical element has its own chemical identity and behaves differently from every other element. Atoms of the same element are identical in chemical behavior; they differ from the atoms of other elements in chemical behavior, in structure, and (usually) in weight.

Atoms and Molecules

A molecule is usually defined as a structure consisting of two or more atoms bound together. A molecule may consist of two or more atoms of the same element—as in hydrogen, sulfur, and phosphorus molecules, for example. On the other hand, a molecule may be made up of atoms of two or more different elements, in which case it constitutes the smallest subdivision of a chemical compound, such as water. Some molecules, such as those of certain proteins, contain many thousands of atoms.

A molecule is sometimes defined—especially by chemists—as the smallest part of an element or compound that can exist alone and still retain the properties of the element or compound. By this definition, the atoms of certain elements (for example, helium and neon) are molecules. Such molecules are termed monatomic.

Atomic Theories

In the fifth century B.C., Empedocles, a Greek philosopher, developed the theory that all matter in the universe was composed of various combinations and proportions of four elementary substances—earth, air, fire, and water. Metals, for example, were considered to be composed of earth and fire, since they could be produced by placing ores (earth) in a flame. The shinier a metal, the more fire it was believed to contain.

"Solid" Atoms

The Greek philosopher Democritus (460?-370 B.C.) accepted the four-element theory of Empedocles and taught that all matter was composed of tiny, indestructible, solid particles (which he called atoms) of earth, air, fire, and water. Democritus believed that spiritual, as well as physical, things were composed of atoms and that atoms could not be subdivided. ("Atom" comes from the Greek word atomos, which means "indivisible.") Democritus was not the first atomist (as persons who subscribed to his views are called), but he had great influence.

The atomism of the early Greeks, although mentioned by such men as Giordano Bruno (1548-1600), Sir Isaac Newton (1642-1727), and Daniel Bernoulli (1700-1782), was largely neglected for more than 2,000 years. John Dalton of England originated the modern atomic theory in the early 19th century. Dalton's description of the atom as a submicroscopic, solid, nondivisible particle of a chemical element seemed to explain satisfactorily the properties of matter that were known at that time.

The concept of solid atoms, although now outmoded, is still useful in diagramming certain chemical changes and in explaining the nature of molecules. For example, ball-like models of various kinds of atoms are joined together to show the structure of large molecules.

Atomic Theory Becomes Complex

In 1897 the English physicist J. J. Thomson determined the nature of electrons—tiny particles with negative electric charges—and measured their mass. He proposed that an atom was made up of a number of electrons embedded in a sphere of positive electricity.

In 1911 the British physicist Ernest Rutherford proposed the nuclear atom. According to his theory, an atom consists of a small, dense, positively charged nucleus (center) with electrons revolving around it. The atom is mostly empty space, because the diameter of the nucleus is much smaller than the total diameter of the atom. In 1913, the Danish physicist Niels Bohr expanded upon this theory. According to Bohr, the electrons circle the nucleus in fixed orbits, and when an atom gains or loses energy the electrons jump from one orbit to another.

Present-day atomic theory is largely based on concepts proposed in the 1920's by the Austrian Erwin Schrdinger, the Germans Werner Heisenberg and Max Born, Paul Dirac of Great Britain, and Louis de Broglie of France. These physicists were pioneers in the field of quantum mechanics.

The atom of quantum mechanics is difficult to picture and is best described mathematically. The negative charge of electrons is treated as a type of wave, to yield equations whose solutions give the probability of finding an electron in various regions around the nucleus. An electron's position and velocity cannot both be known accurately at the same time. The electron is sometimes pictured as a negatively charged cloud covering a large area outside the nucleus; in the region where the probability of finding the electron is greatest, the cloud is densest (and the concentration of charge is the highest).

The Basic Units of the Atom

Three kinds of particles make up atoms: the electron, mentioned in the previous section of this article, the proton, and the neutron. They are commonly referred to as subatomic particles.

The Proton

carries a single positive charge of electricity equal in magnitude to the negative charge of the electron.

The Neutron

carries no electric charge and is therefore electrically neutral. Its mass is approximately equal to that of the proton. A single, isolated neutron is an unstable particle—that is, it rapidly decays into other particles.

The Electron

has a mass much less than that of the proton or neutron (about 1/1836 as great). The electron carries a single negative charge of electricity.

Protons and neutrons make up the nucleus of an atom. For this reason, they are sometimes referred to as nucleons. The nucleus of the simplest atom, ordinary hydrogen, consists of a single proton. In all other atoms, the nucleus is made up of both protons and neutrons.

As the result of many kinds of experiments with devices called particle accelerators, physicists believe that protons and neutrons are themselves composed of smaller particles. These particles are called quarks. Electrons, unlike protons and neutrons, are elementary particles—that is, they are not composed of smaller particles.

Each type of particle has corresponding to it an antiparticle of the same mass but of opposite magnetic and electrical properties. Physicists believe antiprotons, antineutrons, and positrons (antielectrons) can form atoms of antimatter.

Particle Masses

By using devices such as particle accelerators, bubble chambers, and mass spectrographs, physicists have obtained precise measurements of particle masses.

The mass of a proton, neutron, or electron is often expressed in terms of atomic weight, for which the carbon 12 atom (a carbon atom that has 6 protons and 6 neutrons) serves as the standard. The carbon 12 atom is assigned a mass of 12. A free proton (one that is not bound with one or more other particles in an atomic nucleus) has a mass of 1.0072766; a free neutron, a mass of 1.0086654; and an electron, a mass of 0.000548597. The actual mass of the electron is about 9.1 x 10-31 kilogram, which means that it would take 1,100,000,000,000,000,000,000,000,000,000 electrons to weigh one kilogram.

When protons and neutrons combine to form a nucleus, their combined mass is slightly less than the total sum of their individual masses.

Properties of Atoms

Atomic Number

The atomic number is the number of protons contained in an atom's nucleus. When an atom is in its normal state (that is, when it is not ionized), the number of its planetary electrons equals the atomic number. The positive charge of the nucleus is thus equal to the negative charge of the planetary electrons, so that the atom as a whole is electrically neutral. All atoms of the same element have the same atomic number.

Atomic Weight

The atomic weight of an atom is a relative figure that is compared to the standard figure assigned to the carbon 12 atom, and is approximately equal to the number of protons and neutrons in the atom's nucleus. Chemists and physicists sometimes use the term atomic mass when speaking of specific atoms, preferring to use atomic weight to mean the average weight of all the atoms of an element. (As explained later under Isotopes, atoms of the same element can have different numbers of neutrons and, therefore, different weights.)

There are several methods of determining atomic weights. One of the oldest (and most accurate) methods is based on the long-established fact that when chemical elements unite to form compounds, they do so in fixed proportions. For example, careful weighing has shown that when silver reacts with chlorine to form silver chloride, the two elements always combine in the ratio of 0.32867 gram of chlorine for each gram of silver. Since a silver chloride molecule consists of one atom of chlorine and one atom of silver, the ratio of the atomic weights of chlorine and silver must be 0.32867 gram 1 gram, or 0.32867. Therefore, if the atomic weight of either chlorine or silver is known (from a weight ratio relating either of these elements to carbon, for example), the atomic weight of the other element can be readily calculated.

The mass spectrometer, an instrument in which beams of ions (electrically charged atoms or molecules) are deflected by a magnetic field, is also used to determine atomic weights. The masses of the ions in a beam can be calculated from the amount of deflection.

Isotopes

Most chemical elements, as found in nature, are made up of a mixture of atoms of different atomic weights. Atoms with the same atomic number but different atomic weights are called isotopes. Lithium, for example, has two isotopes: lithium 6, with an atomic weight of 6.015; and lithium 7, with an atomic weight of 7.016. Since there are about 12 times as many atoms of lithium 7 found in nature as atoms of lithium 6, the atomic weight of the element lithium is calculated to be 6.939.

The isotope carbon 12, used as the reference standard for the atomic-weight scale, constitutes almost 99 per cent of the carbon found in nature. (Carbon 13, a stable isotope, and carbon 14, a radioactive isotope, constitute the remainder.) Some elements, such as aluminum, phosphorus, and iodine, possess no natural isotopes; they are made up of only one kind of atom.

Isobars

Atoms that have the same approximate atomic weight but different atomic numbers are called isobars. Neptunium 239 and plutonium 239, for example, are isobars—both have 239 nucleons but in neptunium 93 are protons and in plutonium 94 are protons.

Electron Shells

The electronic structure of atoms can be given in terms of a shell model. Each shell is a group of electrons all roughly the same distance from the nucleus. The number of electrons per shell—as well as the total number of electrons—is the same for all the atoms of a given element, provided that each atom is in its normal state. In the atoms now known, the number of electron shells ranges from one to seven. These shells are designated by the letters K, L, M, N, O, P, and Q, in order of their distance from the nucleus.

An atom of chlorine, for example, has 17 electrons in three shells—2 in the K, or innermost, shell; 8 in the L, or second, shell; and 7 in the M shell. A potassium atom, with 19 electrons, has four shells, with the electrons arranged 2, 8, 8, 1.

The chemical properties of the various elements are determined primarily by the number of electrons in the outermost shells of their atoms. For example, the atoms of the metals lithium, sodium, and potassium, which have similar properties, each have a single electron in their outermost shell. It is for this reason that these elements are grouped together in the Periodic Table, a chart that arranges elements according to their properties.

Valence

When atoms take part in a chemical reaction, the electrons in their outermost shells are rearranged. An atom with a single electron in this shell tends to lose the electron. An atom with many electrons in its outermost shell (such as chlorine, with its 7), tends to capture one or more additional electrons until it has a number in this shell that represents a stable arrangement. Groups of 2, 8, 18, and 32 electrons are known to represent highly stable arrangements, and these numbers represent the maximum number of electrons that can exist in the K, L, M, and N shells, respectively, according to quantum mechanics.

Valence is the ability of an atom to gain or lose electrons, or to share them with another atom. An atom displays its valence when it combines with an atom of another element in a chemical reaction.

Chemical and Nuclear Energy

In a chemical reaction, only the planetary electrons of an atom are affected. The nucleus does not take part. In the formation of sodium chloride (common salt) from metallic sodium and gaseous chlorine, for example, the single electron in the outermost shell of the sodium atom is transferred to the outermost shell of the chlorine atom. Chemical energy in the form of heat is liberated, and the sodium and chlorine atoms join together.

Chemical reactions produce heat, light, and electrical energy. Nuclear energy can be released only through bombardment by subatomic particles, through the application of extremely high temperatures, or (in small amounts) by radioactivity.

For books about the atom, See book list at the end of Nuclear Energy.