Electron Energy Levels In The Atom example essay topic

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 for cesare 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 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 atoms, 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 PierreGASSENDI, Robert BOYLE, and especially Isaac NEWTON.

The latter part of Newton's book Optics 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 ages 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 amass 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. Fromthe 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 electromagnetic ism, 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/2 pi), where 'h' is a quantity known asPLANCK'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's tate. 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 ind 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 not wo 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 1920's. 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 1920's 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 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 we restudied intensively in the 1930's and 1940's, 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 breakup 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 FritzSTRASSMANN. 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 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 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 Mellen, Andrew, From Atoms to Atom (1952);
Whittaker, Edmund, A History of the Theories of Aether and Electricity (1960).