Atom and Atomic Theory

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Atom and Atomic Theory, the study of the nature of atoms and the forces which hold them together.
In ancient Greek philosophy the word atom was used to describe the smallest bit of matter that
could be conceived. This “fundamental particle,” to use the present-day term for this concept,
was thought of as indestructible; in fact, the Greek word for atom means “not divisible.” Knowledge
about the size and nature of the atom grew slowly throughout the centuries when people were content
merely to speculate about it.
With the advent of experimental science in the 16th and 17th centuries (see CHEMISTRY; SCIENCE),
progress in atomic theory quickened. Chemists soon recognized that all liquids, gases, and solids can
be analyzed into their ultimate components, or elements (see ELEMENTS, CHEMICAL). For example, salt was
found to be composed of two distinct and different elements, sodium and chlorine, which are joined
together in an intimate form known as a chemical compound. Air was discovered to consist of a mixture of
hydrogen for every atom of oxygen.

Dalton's Theory
John Dalton, a British schoolmaster and chemist, was fascinated by the patchwork puzzle of the elements.
Early in the 19th century he made studies of the way in which the various elements combine with one another
to form chemical compounds. Other scientists, among them the English physicist Sir Isaac Newton, had already
speculated that the smallest units of a substance are atoms. Dalton was regarded as the founder of atomic
theory because he made the theory quantitative. He showed how these atoms link together in definite
proportions. Subsequent investigations proved that the smallest unit of a chemical substance such as water is
a molecule. Each molecule of water consists of a single atom of oxygen and two atoms of hydrogen joined by an
electrical force called a “chemical bond.” See CHEMICAL REACTION.
All atoms of any given element behave in the same way chemically. Thus, from a chemical viewpoint, the atom is
the smallest entity to be considered. The chemical properties of the various elements are quite different; their
atoms combine in many different ways to form a multitude of different chemical compounds. Some elements, such as
the gases helium and argon, are inert, that is, they fail to react with other elements. Unlike oxygen, which has
a diatomic molecule (two atoms combined in a single molecule), helium and other inert gases are monatomic elements,
with a single atom per molecule. See NOBLE GASES.

Avogadro's Law
The study of gases attracted the attention of the Italian physicist Amedeo Avogadro, who in 1811 formulated an
important law bearing his name (see AVOGADRO'S LAW). This law states that equal volumes of different gases contain
the same number of molecules when compared under the same conditions of temperature and pressure. Given these
conditions, two identical bottles, one filled with oxygen and the other with helium, will contain exactly the same
number of molecules. Twice as many atoms of oxygen will be present, however, because oxygen is diatomic.

Atomic Weight
Measurement of the weights of standard volumes (that is, the densities) of different gases permits direct comparison
of the weights of individual gas molecules. When oxygen is taken as a standard and the oxygen atom is assigned a value
of 16.0000 atomic mass units (amu), helium is found to have an atomic weight of 4.003 amu, fluorine 19.000, and sodium
22.997. (Note that it is customary to speak of “atomic weights,” although “atomic masses” would perhaps be more accurate.
Mass is a measure of the quantity of matter in a body, whereas weight is the force exerted on the body by the influence
of gravity. Thus, “atomic weight” is measured in amu. In processes that occur within the nuclei of atoms, such as nuclear
fission, mass is converted into energy.)
The observation that many atomic weights are close to whole numbers led the British chemist William Prout to suggest in
1816 that all elements might be composed of hydrogen atoms. Subsequent measurements of atomic weights revealed that chlorine,
for example, has an atomic weight of 35.455. The discovery of such fractional atomic weights appeared to invalidate Prout's
hypothesis until a century later, when it was discovered that the atoms of most elements do not all have the same
weight. Atoms of the same element that differ in weight are known as isotopes (see ISOTOPE). In the case of chlorine two
isotopes occur in nature. Experiments show that chlorine is a mixture of three parts of chlorine-35 for every one part of
the heavier chlorine-37 isotope. This proportion accounts for the observed atomic weight of chlorine. Atomic scientists can
measure isotopes with great precision. For example, the light isotope of chlorine is measured at 34.97867 amu.
The standard used for the calculation of atomic weights has recently been changed. During the first part of the 20th century
it was customary to use natural oxygen as the standard against which atomic weights or masses were computed; oxygen was
assigned an integral atomic weight of 16. This standard was used by chemists even after the rare isotopes of oxygen
(oxygen-17 and oxygen-18) were discovered in 1929, because the small amounts of these isotopes in natural oxygen are
relatively, although not absolutely, in constant proportion to the abundant isotope, oxygen-16. Physicists found it easier,
however, to compute atomic masses against only the oxygen-16 isotope. This method resulted in two slightly different tables
of atomic weights or masses. The situation was resolved in the early 1960s, when the international unions of chemistry and
physics agreed on a single new standard, the abundant isotope of carbon, carbon-12. The new standard completely replaced the
two earlier standards for all scientists. The new standard is particularly appropriate because carbon-12 is often used as a
reference standard in computations of atomic masses using the mass spectrometer. Moreover, the table of atomic weights based
on carbon-12 is in close agreement with the old table based on natural oxygen.

Periodic Table
By the middle of the 19th century several chemists recognized that similarities in the chemical properties
of various elements implied a regularity that might be illustrated by arranging the elements in a tabular or periodic form.
The Russian chemist Dmitry Mendeleyev proposed a chart of elements called the periodic table (see PERIODIC LAW), in which
the elements are arranged in rows and columns so that elements with similar chemical properties are grouped together.
According to this arrangement, each element was assigned a number (atomic number) ranging from 1 for hydrogen to 92 for
uranium. Because not all the elements were known at the time of Mendeleyev, blank spaces were left in the periodic table,
each of which corresponded to a missing element. Further research, aided by the arrangement of the known elements in the
chart, led to the discovery of missing elements. Elements of higher atomic number have correspondingly heavier atomic weights;
this fact could have been predicted from Prout's hypothesis.

Size of the Atom
Curiosity about the size of the atom and its weight tantalized hundreds of scientists for a long period during which lack of
adequate instruments and proper techniques prevented them from obtaining satisfactory answers. Subsequently, a variety of
ingenious experiments was devised to determine the size and weight of the various atoms. The lightest of all atoms, hydrogen,
has a diameter of 1 × 10-8 cm (0.00000001 cm) and weighs 1.7 × 10-24 (the fraction of a gram represented by 17 preceded by
23 zeros and a decimal point). An atom is so small that a single drop of water contains more than a million million billion
atoms.

Radioactivity
That the atom is not a solid bit of matter, incapable of further subdivision, became evident with the discovery of radioactivity.
In 1896 the French physicist Antoine Henri Becquerel found that certain substances, such as uranium salts, give off penetrating
rays of mysterious origin. Only a year earlier the German scientist Wilhelm Conrad Roentgen had announced the discovery of X
rays, which can penetrate sheets of lead. The French scientists Marie Curie and her husband Pierre Curie contributed further
to an understanding of radioactive substances (see RADIUM). As a result of the research of the British physicist Ernest Rutherford
and his contemporaries, it was shown that uranium and some other heavy elements, such as thorium and radium, emit three different
kinds of radiation, initially called alpha (a), beta (b), and gamma (g) rays. The first two, which were found to consist of
electrically charged bits of matter, are now called alpha and beta particles. Gamma rays eventually were identified as
electromagnetic waves, similar to X rays but of shorter wavelengths (see ELECTROMAGNETIC RADIATION).

Rutherford Nuclear Atom
Recognition of the nature of radioactive emissions enabled physicists to penetrate into the mystery of the atom. Far from
being a solid bit of matter, the atom was found to consist mostly of space. At the center of this space is an infinitesimally
small core called the nucleus. Rutherford established that the mass of the atom is concentrated in its nucleus. He also
proposed that satellites called electrons travel in orbits around the nucleus (see ELECTRON). The nucleus has a positive
charge of electricity; the electrons each have a negative charge. The charges carried by the electrons add up to the same
amount of electricity as resides in the nucleus, and thus the normal electrical state of the atom is neutral.

Bohr Atom
To explain the structure of the atom, the Danish physicist Niels Bohr developed in 1913 a hypothesis known as the Bohr theory
of the atom (see QUANTUM THEORY). He assumed that electrons are arranged in definite shells, or quantum levels, at a
considerable distance from the nucleus. The arrangement of these electrons is called the electron configuration. The number
of such electrons equals the atomic number of the atom; hydrogen has a single orbital electron, helium has 2, and uranium has
92. The electron shells are built up in a regular fashion from a first shell to a total of seven shells, each of which has an
upper limit to the number of electrons that it can accommodate. The first shell is complete with two electrons, the second can
hold up to eight electrons, and successive shells hold still larger numbers. The “last” electrons, those which are outermost or
added last to the atom's structure, determine the chemical behavior of the atom.
The inert, or noble, gases (helium, neon, argon, krypton, xenon, and radon) all have completely filled outer shells. They do
not enter into chemical combinations in nature, although the three heaviest inert gases (krypton, xenon, and radon) have formed
chemical compounds in the laboratory. On the other hand, the outermost shells of such elements as lithium, sodium, and potassium
contain only one electron. These elements combine readily with other elements (transferring their outermost electrons to them)
to form a great many chemical compounds.
Atomic shells do not necessarily fill up with electrons in consecutive order. The electrons of the first 18 elements in the
periodic table are added in a regular manner, each shell being filled to a designated limit before a new shell is started.
Beginning with the 19th element, the outermost electron starts a new shell before the previous shell is completely filled.
A regularity is still maintained, however, as electrons fill successive shells in a repetitious back-and-forth pattern.
The result is the regular repetition of chemical properties for atoms of increasing atomic weight that corresponds to the
arrangement of the elements in the periodic table.
It is convenient to visualize the electrons moving about the nucleus of an atom much as if they were planets moving
about the sun. This view is much more precise than that held by contemporary physicists, however. It is now known that
it is impossible to pinpoint the precise position of an electron in the atom's space without disturbing its predicted location
at some future time. This uncertainty is resolved by attributing to the atom a cloudlike form, in which the electron's
position is defined in terms of the probability of finding it at some distance from the nucleus. This rather fuzzy schematic
conception of the atom may be reconciled with the solar-system model by noting that in the tiny space of the atom the
electron, which makes many billions of orbits around the nucleus in a single second, is everywhere at once. The cloud view
thus gives a form to the atom that is not supplied by a solar-system model.

Line Spectra
One of the great successes of theoretical physics was the explanation of the characteristic line spectra of various elements
(see SPECTROSCOPY: SPECTRUM LINES). Atoms excited by a supply of energy from an external source emit light of well-defined
frequencies. If hydrogen gas, for example, is held at low pressure in a glass tube and an electrical current is passed through
it, visible light of a reddish color is given off. Careful examination of this light with a prism spectroscope shows a line
spectrum, a series of regularly spaced lines of light, each of which has a definite wavelength and associated energy. The
Bohr theory permits the physicist to calculate these wavelengths in a straightforward fashion. It is assumed that in the
hydrogen atom the outer electron can move in stable orbits. While the electron remains in an orbit at a fixed distance from
the nucleus, the atom does not radiate energy. When the atom is excited, the electron jumps to a higher-energy orbit farther
from the nucleus, and as it falls back to its normal orbit, it emits a discrete amount of energy corresponding to a certain
wavelength of light. Each line of light observed represents an electronic transition between a higher and lower energy orbit.
In many heavier elements, if an atom is sufficiently excited so that inner electrons close to the nucleus are affected, then
penetrating radiation, or X rays, will be emitted. These electronic transitions involve large amounts of energy.

Atomic Nucleus
In 1905 Albert Einstein developed his mass-energy equation, E = mc2, as part of his special theory of relativity. This equation
states that with a given mass (m) is associated an amount of energy (E) equal to this mass multiplied by the square of the
velocity of light (c). A very small amount of mass is equivalent to a vast amount of energy. Because more than 99 percent of
the atom's mass is in the nucleus, any release of the atom's energy would have to come from the nucleus.
In 1919 Rutherford exposed nitrogen gas to a radioactive source that emitted alpha particles. Some of the alpha particles collided
with the nuclei of the nitrogen atoms. As a result of these collisions, the nitrogen atoms were transmuted into oxygen atoms.
A positively charged particle was emitted from the nucleus of each of the atoms undergoing transmutation. These particles were
recognized as being identical to the nuclei of hydrogen atoms. They are called protons (see PROTON). Although further research
proved that protons are constituents of the nuclei of all elements, no more clues to the structure of the nucleus were found
until 1932, when the British physicist Sir James Chadwick discovered in the nucleus another particle, known as the neutron,
having the same weight as the proton but without an electrical charge. It was then realized that the nucleus is made up of
protons and neutrons. In any given atom, the number of protons is equal to the number of electrons and hence to the atomic
number of the atom. Isotopes are then explained as atoms of the same element (that is, containing the same number of protons)
that have different numbers of neutrons. In the case of chlorine, one isotope is identified by the symbol of 35Cl and its
heavy relative by 37Cl. The superscripts identify the mass number of the isotope and are numerically equal to the total number
of neutrons and protons in the nucleus of the atom. Sometimes the atomic number is given as a subscript, as in }Cl.
The least stable arrangement of nuclei is one in which an odd number of neutrons and an odd number of protons are present; all
but four isotopes containing nuclei of this kind are radioactive. The presence of a large excess of neutrons over protons
detracts from the stability of a nucleus; nuclei in all isotopes of elements above bismuth in the periodic table contain this type
of arrangement, and they are all radioactive. Most known stable nuclei contain an even number of protons and an even number of
neutrons.

Artificial Radioactivity
Experiments by the French physicists Frédérick and Irène Joliot-Curie in the early 1930s showed that stable atoms of an element
may be made artificially radioactive by suitable bombardment with nuclear particles or rays. Such radioactive isotopes
(radioisotopes) are produced as a result of a nuclear reaction, or transformation. In such reactions the 270-odd isotopes found
in nature serve as targets for nuclear projectiles. The development of atom smashers, or accelerators, for hurling these
projectile-particles to high energy has made it possible to observe thousands of nuclear reactions.

Nuclear Reactions
In 1932 two British scientists, Sir John D. Cockcroft and Ernest T. S. Walton, were the first to use artificially accelerated
particles to successfully disintegrate the nucleus. They produced a beam of protons, which were boosted to high speed by means
of a high-voltage device called a voltage multiplier. These particles were then used to bombard a lithium target. In this nuclear
reaction, lithium-7 (7Li) splits into two fragments, which are nuclei of helium atoms. The reaction is expressed by the equation


7Li + 1H = 4He + 4He

Physicists have measured the weights of these atoms precisely—7Li has a weight of

7.018242 amu; 1H, 1.008137 amu; and 4He, 4.003910 amu.

The weights on the left side of the equation add up to

8.026379 amu, whereas those on the right side total 8.007820 amu;
a “loss” of 0.018559 amu has occurred.

Using Einstein's E = mc2 relation, 1 amu is found to be the equivalent of 931.3 million electron volts (MeV) of energy. On this
basis the nuclear reaction with lithium releases 17.28 MeV of energy. The “lost” mass appears as energy in the form of the
violent motion of the helium nuclei. See NUCLEAR CHEMISTRY.

Particle Accelerator
The American physicist Ernest O. Lawrence developed about 1930 a particle accelerator called a cyclotron. This machine generates
electrical attractive and repulsive forces that accelerate atomic particles while they are confined to a circular orbit by the
electromagnetic force of a large magnet. The particles spiral outward under the influence of these electric and magnetic forces,
reaching extremely high speeds. The acceleration takes place in a vacuum so that the particles do not collide with molecules
of air. Because the equipment necessary for producing intense magnetic forces is massive, high-energy machines are huge and
expensive installations. See PARTICLE ACCELERATORS.

Nuclear Forces
Modern nuclear theory is based on the notion that nuclei consist of neutrons and protons that are held together by extremely
powerful “nuclear” forces. The elucidation of these nuclear forces requires physicists to disrupt neutrons and protons by
bombarding nuclei with extremely energetic particles. Such bombardments have revealed more than 200 so-called elementary particles,
or tiny bits of matter, most of which exist for much less than one hundred-millionth of a second.
This subnuclear world was first revealed in cosmic rays (see COSMIC RAYS). These rays consist of highly energetic particles
that constantly bombard the earth from outer space, penetrating down through the atmosphere and even into the earth's crust.
Cosmic radiation includes many types of particles, some having energies far exceeding anything achieved in particle accelerators.
When these energetic particles strike nuclei, new particles are created. Among the first such particles to be observed were the
muons (detected in 1937) and pions (1947). The existence of the pion had been predicted in 1935 by the Japanese physicist Yukawa
Hideki. According to the most widely accepted theory, nuclear particles are held together by “exchange forces,” in which pions
common to both neutrons and protons are continuously exchanged between them. The binding of protons and neutrons by pions is similar
to the binding of two atoms in a molecule through sharing or exchanging a common pair of electrons. These particles are about 200
times as heavy as electrons. The muon is essentially a heavy electron and can be either positively or negatively charged. The pion,
slightly heavier than the muon, can carry a positive or negative charge, or no charge.

Elementary Particles
Accelerator studies eventually established that each kind of particle also has an antiparticle of the same mass but opposite
in charge or other electromagnetic property. Physicists have long sought a theory that would put this bewildering array of
particles in order. Particles are now grouped according to the force that usually controls their interactions. Hadrons
(strong nuclear force) include hyperons, mesons, and the neutron and proton. Leptons (electromagnetic and weak forces)
include the tau, muon, electron, and neutrinos. Bosons (particlelike objects associated with interactions) include the photon
and the hypothetical carriers of the weak force and of gravitation. The weak nuclear force is evident in such radioactive or
particle-decay reactions as alpha decay (the release of a helium nucleus from an unstable atomic nucleus). See ANTIMATTER.
In 1963 the U.S. physicists Murray Gell-Mann and George Zweig proposed that hadrons are actually combinations of more fundamental
particles called quarks, the interactions of which are carried by particlelike gluons. This theory underlies current investigations
and has served to predict the existence of further particles. See ELEMENTARY PARTICLES; GLUON; QUARK.

Release of Atomic Energy
Two nuclear processes of great practical significance because they provide vast amounts of energy are fission, the splitting
of a heavy nucleus into lighter ones, and thermonuclear fusion, the fusion of two light nuclei (at extremely high temperatures)
to form a heavier one. The Italian-born American physicist Enrico Fermi achieved fission in 1934, but the reaction was not
recognized as such until 1939, when the German scientists Otto Hahn and Fritz Strassmann announced that they had split uranium
nuclei by bombarding them with neutrons. Neutrons are also released by the reaction and can cause a chain reaction with other
nuclei. An uncontrolled chain reaction is seen in the explosion of an atomic bomb. Heat from controlled reactions, however,
as in nuclear reactors, can be used to produce electric power. Thermonuclear fusion occurs in stars, including the sun,
and is the source of their heat and light. Uncontrolled fusion is seen in the explosion of a hydrogen bomb, but physicists
are currently trying to develop a practical controlled-fusion device. See NUCLEAR ENERGY; NUCLEAR WEAPONS.

Contributed by:
Chen Ning Yang
Alvin M. Weinberg