Essay/Term paper: The atom

Essay, term paper, research paper:  Chemistry

An atom is the smallest unit of matter that is
recognizable as a chemical ELEMENT. Atoms of
different elements may also combine into systems
called MOLECULES, which are the smallest units
of chemical COMPOUNDS. In all these ordinary
processes, atoms may be considered as the
ancient Greeks imagined them to be: the ultimate
building blocks of matter. When stronger forces
are applied to atoms, however, the atoms may
break up into smaller parts. Thus atoms are
actually composites and not units, and have a
complex inner structure of their own. By studying
the processes in which atoms break up, scientists
in the 20th century have come to understand many
details of the inner structure of atoms. The size of
a typical atom is only about 10 (-10th) meters. A
cubic centimeter of solid matter contains
something like 10 (24th) atoms. Atoms cannot be
seen using optical microscopes, because they are
much smaller than the wavelengths of visible light.
By using more advanced imaging techniques such
as electron microscopes, scanning tunneling
microscopes, and atomic force microscopes,
however, scientists have been able to produce
images in which the sites of individual atoms can
be identified. EARLY ATOMIC THEORIES The
first recorded speculations that MATTER
consisted of atoms are found in the works of the
Greek philosophers LEUCIPPUS and
DEMOCRITUS. The essence of their views is
that all phenomena are to be understood in terms
of the motions, through empty space, of a large
number of tiny and indivisible bodies. (The name
"atom" comes from the Greek words atomos, for
"indivisible.") According to Democritus, these
bodies differ from one another in shape and size,
and the observed variety of substances derives
from these differences in the atoms composing
them. Greek atomic theory was not an attempt to
account for specific details of physical phenomena.
It was instead a philosophical response to the
question of how change can occur in nature. Little
effort was made to make atomic theory
quantitative--that is, to develop it as a scientific
hypothesis for the study of matter. Greek atomism,
however, did introduce the valuable concept that
the nature of everyday things was to be
understood in terms of an invisible substructure of
objects with unfamiliar properties. Democritus
stated this especially clearly in one of the few
sayings of his that has been preserved: "Color
exists by convention, sweet by convention, bitter
by convention, in reality nothing exists but atoms
and the void." Although adopted and extended by
such later ancient thinkers as EPICURUS and
LUCRETIUS, Greek atomic theory had strong
competition from other views of the nature of
matter. One such view was the four-element
theory of EMPEDOCLES. These alternative
views, championed by ARISTOTLE among
others, were also motivated more by a desire to
answer philosophical questions than by a wish to
explain scientific phenomena. ORIGINS OF
MODERN ATOMISM When interest in science
revived in Europe in the 16th and 17th centuries,
enough was known about Greek atomism to form
the basis for further thought. Among those who
revived the atomic theory were Pierre
GASSENDI, Robert BOYLE, and especially
Isaac NEWTON. The latter part of Newton's
book Optiks is a series of detailed speculations on
the atomic nature of matter and light, indicating
how some of matter's properties are to be
understood in terms of atoms. In the 19th century,
two independent lines of reasoning strengthened
the belief of most scientists, by then, in the atomic
theory. Both approaches also began to reveal
some quantitative properties of atoms. One
approach, pioneered by John DALTON, involved
chemical phenomena. The other, involving the
behavior of gases, was carried out by physicists
such as Rudolph CLAUSIUS and James Clerk
MAXWELL. Dalton's main step forward was his
introduction of ATOMIC WEIGHTS. Dalton
studied the elements then known and analyzed the
data of their reactions with one another. He
discovered the law of multiple proportions, which
states that when several distinct reactions take
place among the same elements, the quantities that
enter the reactions are always in the proportions of
simple integers--that is, 1 to 1, 2 to 1, 2 to 3, and
so on. From this came the concept that such
reacting quantities contain equal numbers of atoms
and are therefore proportional to the masses of
individual atoms. Dalton gave the lightest known
element, hydrogen, an atomic weight of 1, and
developed comparative atomic weights for the
other known elements accordingly. The study of
gases in terms of atomic theory was begun by
Daniel BERNOULLI in the 18th century.
Bernoulli showed that the pressure exerted by a
gas came about as the result of collisions of the
atoms of the gas with the walls of its container. In
1811, Amadeo AVOGADRO suggested that
equal volumes of different gases, under the same
conditions of pressure and temperature, contain
equal numbers of atoms. The number of atoms in a
mass of gas equal to one gram atomic weight--a
quantity of an element, in grams, that has the same
numerical value as the element's atomic weight--is
now known to be approximately 6.022 x 10
(23rd). This huge value is an indication of the
disparity in size between atoms and everyday
objects. Avogadro himself never estimated the
magnitude of this value, although it is now known
as the AVOGADRO NUMBER. Estimates of its
value were first given in the mid-19th century by
Clausius and Maxwell. An accurate measurement
was not carried out until the early 20th century,
using the diffraction of X-rays by crystals. From
the value of the Avogadro number it is possible to
infer the mass of the individual atoms, which for
hydrogen turns out to be 1.6 x 10 (-24th) grams.
DISCOVERY OF THE ELECTRON AND OF
RADIATION By the end of the 19th century
almost all scientists had become convinced of the
truth of the atomic theory. By that time, ironically,
evidence was just beginning to accumulate that
atoms are not in fact the indivisible particles
suggested by their name. One source of such
evidence came from studies using gas discharge
tubes, which are similar to neon lights. In such
tubes, a gas at low pressure is subjected to intense
electrical forces. Under these conditions, various
colored glows (now known as glow
DISCHARGE) are observed to traverse the tube.
One blue glow at one end of the tube, around the
electrode known as the CATHODE, was
observed for a wide variety of gases. The glow
was shown by Joseph John THOMSON in 1897
to involve a stream of negatively charged particles
with individual masses much smaller than that of
any atom. These particles were called
ELECTRONS, and they were soon recognized to
be a constituent of all atoms. That is, atoms are
not indivisible but contain parts. In the late 19th
and the early 20th century it was also found that
some kinds of atoms are not stable. Instead they
transform spontaneously into other kinds of atoms.
For example, uranium atoms slowly change into
lighter thorium atoms, which themselves change
into still lighter atoms, eventually ending up as
stable atoms of lead. These transformations, first
observed by Antoine Henri BECQUEREL, came
to be known as RADIOACTIVITY, because the
atomic changes were accompanied by the
emission of several types of radiation. Atoms are
ordinarily electrically neutral. Therefore the
negative charge of the electrons in an atom must
be balanced by a corresponding positive charge.
Because the electrons have so little mass, the
positive constituents of an atom must also carry
most of the atom's mass. The obvious question
arose as to how these varied parts are arranged
within an atom. The question was answered in
1911 through the work of Ernest RUTHERFORD
and his collaborators. In their experiments they
passed alpha particles--a type of radiation emitted
in some radioactive decays--through thin gold
foils. They observed that in some instances the
alpha particles emerged in the opposite direction
from their initial path. This suggested a collision
with a heavy object within the atoms of the gold.
Because electrons are not massive enough to
produce such large deflections, the positive
charges must be involved. Analyzing the data,
Rutherford showed that the positive charge in an
atom must be concentrated in a very small volume
with a radius less than 10(-14th) meter, or one
ten-thousandth the size of the whole atom. This
part of the atom was soon called the nucleus.
Later measurements showed that the size of a
nucleus is approximately given by multiplying the
cube root of the atomic weight by 10(-15th)
meter. RUTHERFORD MODEL Rutherford
proposed an atomic model in which the atom was
held together by electrical attraction between the
nucleus and the electrons. In this model the
electrons traveled in relatively distant orbits around
the nucleus. The model eventually proved
successful in explaining most of the phenomena of
chemistry and everyday physics. Subsequent
studies of the atom divided into investigations of
the electronic parts of the atom, which came to be
known as atomic physics, and investigations of the
nucleus itself, which came to be known as nuclear
physics. This division was natural, because of the
immense difference in size between the nucleus
and the electron orbits and the much greater
energy needed to produce nuclear as compared to
electronic changes. The Rutherford model of the
atom, however, had to face two immediate
problems. One was to account for the fact that
different atoms of the same element behaved in
physically and chemically similar ways. According
to the Rutherford model, electrons could move in
any of the infinite number of orbits allowed by
Newtonian physics. If that were so, different
atoms of the same element could behave quite
differently. (This is actually a problem for any
atomic model based on Newtonian physics, and it
had already been recognized by Maxwell in
1870.) The other problem was that, according to
the principles of electromagneticism, electrons
should continuously emit radiation as they orbit in
an atom. This would cause the electrons to lose
energy and to spiral into the nucleus. It was
estimated that for the single electron in a hydrogen
atom, this would take place in 10(-9th) seconds.
In reality, hydrogen atoms are indefinitely stable.
An important step toward solving these problems
was taken by Niels BOHR in 1913. According to
Bohr, the electrons in atoms cannot exist in
arbitrary orbits. Instead they are found only in
certain "states". The states in which they can exist
are those in which the ANGULAR
MOMENTUM of their orbits is an integer multiple
of h/2pi), where "h" is a quantity known as
PLANCK'S CONSTANT. This constant had
been introduced by Max PLANCK in his theory
describing BLACKBODY RADIATION. BOHR
MODEL According to the Bohr model of the
atom, there is a so-called ground state for any
atom. This ground state has the lowest energy
allowed to the atom, and it is the same for all
atoms containing the same number of electrons.
An atom normally exists in this ground state, which
determined the observed properties of a given
element. Furthermore, according to Bohr, no
radiation is emitted by an atom in its ground state.
This is because energy must be conserved in the
radiation process, and no available state of lower
energy exists for the atom to balance any energy
lost through radiation. An atom can be removed
from its ground state only when enough energy is
given to it, by radiation or collisions, to raise an
electron to an "excited" state. For most atoms this
excitation energy corresponds to several
ELECTRON VOLTS. When the atom is excited,
it will usually emit electromagnetic radiation rapidly
and return to the ground state. The radiation is
emitted in the form of individual packets or quanta,
of light, called PHOTONS. Each photon has an
energy equal to the difference between the energy
of the excited states and the ground state of the
atom. According to a formula developed by
Planck and Albert EINSTEIN, this energy
corresponds to a specific wavelength of the
emitted light. Using his assumption about the
allowed angular momenta for electrons, Bohr was
able to calculate the precise wavelengths in the
SPECTRUM of the simplest atom, hydrogen. The
agreement of his results with observations did
much to convince scientists of the accuracy of his
model. ATOMIC PHYSICS AND QUANTUM
THEORY Bohr was able to extend his atomic
theory to describe, qualitatively, the chemical
properties of all the elements. Each electron in an
atom is assigned a set of four so-called quantum
numbers. (These numbers correspond to the
properties of energy, total orbital angular
momentum, projection of orbital angular
momentum, and projection of spin angular
momentum.) It is also assumed--as had first been
suggested by Wolfgang PAULI in 1924--that no
two electrons in an atom can have the same values
for all four quantum numbers. This came to be
known as the EXCLUSION PRINCIPLE. This
principle influences the way in which the chemical
properties of an element depend on its ATOMIC
NUMBER (the number of electrons in each atom
of the element). A maximum number of electrons
can occur for each energy level, and no more than
that. For example, the lowest energy level of an
atom--the one in which the electrons have zero
orbital angular momentum--can contain up to two
electrons. The one electron in a hydrogen atom
exists at this energy level, as do the two electrons
in a helium atom. For the next heavier atom,
lithium, one of its three electrons must exist in a
higher energy state, and as a result this electron
can more easily be lost to another atom. Those
electrons with approximately the same energy are
said to form a "shell." When an atom contains the
maximum number allowed for some energy level,
that shell is said to be closed. Atoms of INERT
GASES such as helium and argon have all their
shells closed. Although Bohr's model gives a
qualitatively accurate description of atoms, it does
not give quantitatively accurate accurate results for
atoms more complex than hydrogen. In order to
describe such atoms, it is necessary to use
QUANTUM MECHANICS. This theory of
atomic subatomic phenomena was created by
Erwin SCHRODINGER, Werner
HEISENBERG, Paul DIRAC, and others in the
1920s. In quantum mechanics, the electron orbits
are replaced by PROBABILITY distributions that
only indicate in which regions of space each
electron is most likely to be found. An equation
first written by Schrodinger allows this distribution
to be calculated for each atom. From the
distribution, properties of the atom such as energy
and angular momentum can be determined.
Calculations of a wide variety of atomic
phenomena have been carried out by means of
quantum mechanics. Without exception, these
calculations have proven to give an accurate
description of the properties and behavior of
atoms. For the simplest atoms, the observations
and calculations sometimes agree to better than
one part in a billion. EXPLORATION OF THE
NUCLEUS As described above, physicists by the
late 1920s were convinced that they sufficiently
understood the electronic structure of atoms.
Attention therefore turned to the nucleus. It was
already known that nuclei sometimes change into
one another through radioactive decay. Rutherford
had also shown, in 1919, that this could be
accomplished artificially by bombarding nitrogen
nuclei with high-energy alpha particles. In the
process the nitrogen nucleus is converted into an
oxygen nucleus, and a hydrogen nucleus, or
PROTON, is ejected. It had further been
discovered by Thomson, Francis William
ASTON, and others that for a given element the
nucleus sometimes occurs in several different
forms that differ in mass. These chemically similar
but physically distinct atoms were called
ISOTOPES. All of this provided evidence that
atomic nuclei also had some kind of internal
structure that could be explored through
experiments and calculations. Differences in the
integer values of the electric charge and of the
mass of many nuclei soon indicated that protons
were not the only kind of particle to be found
there. That is, the electric charge of a nucleus is
always exactly an integer multiple of the charge of
a proton, so knowledge of this electric charge
always indicates how many protons a nucleus
contains. The mass of a nucleus is also
approximately--but not exactly--an integer
multiple of the mass of a proton. For many atoms,
however, these two integer values are not the
same. For example, a helium nucleus has twice the
charge but four times the mass of a proton.
Clearly, nuclei contain something other than
protons. This problem was solved in 1932 with
the discovery by James CHADWICK of the
NEUTRON. This is a particle that has no electric
charge and is slightly more massive than a proton.
Thus most nuclei are composed of both protons
and neutrons, which collectively are known as
nucleons. A helium nucleus contains two protons
and two neutrons, which correctly give the total
charge and mass of the nucleus. The isotopes of
any given element contain equal numbers of
protons but different numbers of neutrons. For
example, an isotope of hydrogen called
DEUTERIUM contains one proton and one
neutron, and a heavier isotope called TRITIUM
contains one proton and two neutrons. The
problem then arose as to how atomic particles
could be held together in such a small region as the
nucleus. The force holding them had to be different
from others then known to physicists. It was
stronger than the electric forces that can break
electrons away from nuclei. On the other hand, the
nuclear forces between different nuclei that are far
apart are very weak, much weaker than electric
forces at such distances. Nuclear forces were
studied intensively in the 1930s and 1940s, and
many details about their properties were learned.
Ultimately, such studies became a part of the study
of FUNDAMENTAL PARTICLES. NUCLEAR
FORCES AND REACTIONS Measurements of
nuclear masses showed that the mass of a nucleus
is not exactly the sum of the masses of its
constituents. Instead, the total mass is slightly
smaller than this sum. The force binding nuclear
particles together--the so-called BINDING
ENERGY--was linked to this decrease in total
mass. That is, Einstein's equating of mass with
energy indicated that the missing mass constituted
the binding energy required to bring the nuclear
particles together. The stability of a nucleus can be
measured by the magnitude of its binding energy
divided by its number of nucleons. Greater values
for the result correspond to greater stability for a
given nucleus. For lighter nuclei the average
binding energy is small. It tends to increase with
increasing nucleon number, up to nuclei with about
60 nucleons. These are the most stable nuclei.
Beyond that nucleon number, the magnitude of the
average binding energy decreases slowly. The
heaviest known atomic nuclei are the least stable
ones. By comparing the average binding energy of
various nuclei, it is possible to tell whether a
reaction among those nuclei will release energy or
will require extra energy to make it happen.
Reactions between two light nuclei, such as the
combining of two deuterons to produce helium,
generally release helium. Because two nuclei repel
each other electrically, however, such FUSION
occurs only when they are moving fast enough to
overcome this repulsion and can approach one
another to within a short enough distance for the
attraction of the nuclear forces to bring them
together. High-energy fusion reactions are the
source of energy of most stars, and they are also
the means by which all of the elements in the
universe other than hydrogen have been produced.
Very heavy nuclei, on the other hand, can break
up into two or more similar nuclei, liberating
energy in the process. Because of this tendency,
all nuclei containing more than about 210 nucleons
are unstable against various kinds of radioactive
decay. An important example of this instability of
heavy nuclei is nuclear FISSION, discovered in
uranium in 1938 by Otto HAHN and Fritz
STRASSMANN. In fission, the products of the
breakup are two intermediate-sized nuclei and
several neutrons. Fission can happen either
spontaneously or as the result of subjecting the
original nucleus to outside stimulation. The most
important such stimulus is the absorption of a
neutron by the nucleus. Because neutrons are
uncharged, they are not repelled electrically by
nuclei. Thus even very low-energy neutrons can be
absorbed and stimulate fission. In the fission of a
heavy nucleus such as uranium, hundreds of
millions of electron volts of energy are liberated,
millions of times more than in chemical processes
involving the electrons in an atom. Furthermore,
the fact that additional neutrons are liberated in the
fission process allows the possibility of a chain
reaction, in which more and more nuclei are
fissioned as the reaction proceeds. It is such chain
reactions that occur in nuclear-power reactors and
in fission-based nuclear explosives. NUCLEAR
MODELS As a result of studies of nuclear
processes, several models exist to describe the
structure of atomic nuclei. Because neutrons and
protons each satisfy the exclusion principle, this
leads to a shell-structure model of nuclei. In the
so-called independent-particle shell model, each
nucleon is assumed to move under the influence of
an average force produced by the other nucleons.
The energy levels of this motion are described by
quantum mechanics in a way similar to that of
electron energy levels in the atom. This model
helps to explain why certain nuclei, such as the
isotopes of helium that has four nucleons in its
nucleus, have especially high binding energies
compared to nuclei fairly close to them in atomic
weight. Some properties of nuclei, however, are
not well explained by the independent-particle
model. For example, it does not account for the
fact that some nuclei are cigar-shaped rather than
spherical. Other nuclear models have been
proposed to account for such properties.
RECENT WORK IN ATOMIC AND
NUCLEAR PHYSICS Much recent work in
atomic physics has concentrated on atoms in
abnormal situations. For example, studies have
been made of so-called Rydberg atoms, in which
a single electron of a many-electron atom is
excited to a very energetic state. Such Rydberg
atoms behave similarly to hydrogen atoms, and
their properties are accurately described by the
energies calculated from the Bohr theory. There
have also been studies of "exotic" atoms in which
one of the electrons is replaced by a heavier,
negatively charged subatomic particle such as an
antiproton. Because the heavier particle is much
closer to the nucleus than an electron would be,
such atoms serve as a useful probes of nuclear
structure. Nuclear physicists have found methods
for studying nuclei heavier than uranium, which do
not occur naturally. One way to produce
TRANSURANIUM ELEMENTS is by colliding
two beams of lighter nuclei. In such a collision, the
two nuclei sometimes fuse into a heavier nucleus
that can be studied for a short time before it
disintegrates. Such heavy-ion collisions have
produced nuclei that contain as many as 300
nucleons. Gerald Feinberg Bibliography: Beyer,
Robert, ed., Foundations of Nuclear Physics
(1949); Feinberg, Gerald, What is the World
Made Of? (1977); Lapp, Ralph, and Andrews,
Howard, Nuclear Radiation Physics (1972); Pais,
Abraham, Inward Bound (1986); Van Melsen,
Andrew, From Atomos to Atom (1952);
Whittaker, Edmund, A History of the Theories of
Aether and Electricity (1960).