Our modern concept of the atom has not evolved at a uniform rate or by a route that could be called a direct one. Certainly, the atomic theory is not yet complete. The information we now have concerning the structure of the atom is a tribute to generations of scientists who have devoted their lives to the highest standards of scientific endeavor. A study of the lives of many of these scientists would be a fascinating and rewarding experience for any student.

Famous Scientists

The Greek philosopher Democritus (460-370 BC), around 400 B.C., proposed what is considered to be the first atomic theory. For some time, men had wondered whether the smallest pieces of clay and stone would not be just like the smallest pieces of sand and soil. Democritus proposed that the world is composed of two things, void (empty space) and tiny particles which he called "atoms," from the Greek word atomos, meaning indivisible. Atoms, Democritus said, were indestructible, indivisible, and exceedingly small. They could not be broken up because they were the smallest possible particles which could exist. In addition, he proposed that these atoms were many in number and took various shapes. There was a different type of atom for each material in the world; for example, there were atoms of air, atoms of rock, and atoms of iron. Atoms of different materials were about the same size, but each had a different shape. Atoms of solid materials were rough and would not slide easily over one another, while atoms of liquids were smooth. His suggestion was so general and science was so undeveloped at the time, that these ideas could not be tested by experimentation. This early atomic theory by Democritus was based upon intuition and faith that a simple explanation underlay the complexity of everyday observations. The Roman poet Titus Lucretius Carus (c.95-55 BC) voiced his belief in this early atomic theory in his poem De Rerum Natura. Lucretius' poem was, like Democritus' belief, unsupported by experimental evidence. These beliefs concerning the fundamental atomic nature of matter were not widely accepted for centuries to come, because of the domination of the philosophy of Aristotle (384-322 BC). Aristotle and his followers believed that matter was continuous and made up of only one substance called hyle. Aristotle's views were an extension of an idea of another Greek philosopher, Empedocles (c.490-c.430), who said that all matter was made up of fire, earth, air, and water. The influence of the Aristotelian theory was predominant until the seventeenth century, when doubts and objections began to be expressed. Issac Newton (1642-1727) and Robert Boyle (1627-1691) published articles expressing a belief in the atomic nature of elements. Even the philosopher Rene Descartes (1596-1650) believed in a "smallest particle," although he would not admit the existence of void. These works offered not proof; they were explanations of the known, with no predictions of the unknown. It was up to the English chemist, John Dalton (1766-1844), to offer a reasoned hypothesis concerning the existence of atoms.


During the early 1800's, Dalton was puzzled by certain facts which had been observed by experimenters. These facts had to do with chemical changes. Chemical changes were said to occur when one or more materials were transformed to other materials and an energy change accompanied the transformation. The experimenters had worked with many of these chemical changes, such as the digestion of food and the burning of a candle.

One observation based on study of this type was made by Antoine Lavoisier (1743-1794), a French chemist, and was later adopted as a law. Lavoisier had observed that if a chemical change took place in a closed space, the mass of the materials present after the change was equal to the mass of the materials present before the change. The constancy of mass in such transformations was found to be true in all tests of chemical changes. The law was then stated as: MATTER CANNOT BE CREATED NOR DESTROYED. In ordinary chemical changes, then, matter can be altered in many ways; but it cannot be created nor destroyed. Today this law is called the law of conservation of mass, as noted in Section 1.3.


The work of another French chemist, Joseph Proust (1754-1826), also attracted the attention of Dalton. Proust's work, like Lavoisier's, was concerned with mass relations in chemical changes. Proust had observed that particular substances always contain their components in the same proportions by mass. For example, common table salt is made of sodium and chlorine. The ratio of the mass of sodium to the mass of chlorine in any sample of pure salt is always the same. It makes no difference where the sample is obtained, how it is obtained, or how large it is; the mass of sodium will always be in a constant ratio to the mass of chlorine. This principle is known as the law of definite proportions, as noted in Section 2.8.


Dalton was attempting to correlate these findings into a broader generalization when he formulated the beginnings of our present atomic theory. He proposed that all matter is composed of very small particles called atoms, and that these atoms are indivisible. Although, in these two respects, Dalton's hypothesis was like that of Democritus, Dalton believed that atoms were more basic than particles of air or rock, and that atoms of different elementary substances were quite unlike. In Dalton's view, each elementary substance was composed of atoms that were exactly alike, but quite different from atoms of all other elementary substances. He proposed that atoms can unite with other atoms in simple numerical ratios to form compound substances. These last two sentences are the key to the atomic theory. At that time, an element was defined as a material whose atoms are all alike. These elements can react (unite) with each other to form other substances. To use the example of salt again, sodium, a silvery metal, reacts with chlorine, a poisonous, yellowish-green gas, to form the edible, white crystals of common table salt.

We can see how well Dalton's hypothesis explains the two laws he was considering. If atoms are indestructible, then in a chemical change, the atoms must simply be rearranged, and the total number and kind of atoms must remain the same. Therefore, the mass before a reaction must equal the mass after a reaction. Also, if the atoms of an element are always alike, then the mass of each atom of a particular element must equal the mass of every other atom of that element.

In the example of salt, according to Dalton, all sodium atoms equal each other in mass, and all chlorine atoms equal each other in mass. When a sodium atom combines with a chlorine atom, salt is formed. The ratio of the mass of sodium to the mass of chlorine in this pair of atoms is the ratio of the masses of the two atoms. The same is true of any other pair of these atoms. Dalton would say that, since all atoms of sodium are equal in mass and all chlorine atoms are equal in mass, the ratio of the mass of sodium to the mass of chlorine must be the same for any sample of salt. This same type of reasoning would hold true for any particular material. Experimentation has shown that Dalton's hypothesis is not entirely correct. As we will see later, not all atoms of the same element are exactly alike in all respects. However, by taking this fact into consideration, and by changing the word "mass" to "average mass", we can adapt Dalton's hypothesis to present-day knowledge.


All Matter Is Composed Of Atoms
All Atoms Of The Same Element Are Identical
Atoms Of Different Elements Are Different
Atoms Unite In Definite Ratios To Form Compounds

Dalton invented pictorial symbols for the elements and combined these symbols to represent compounds. Note that Dalton believed that the formula for water was HO: Dalton's Symbols & Formulas.


Dalton's atomic theory, which was based on experimental observations, provided an explanation for the law of definite composition and the law of conservation of matter. In addition, the atomic theory was used to predict the LAW OF MULTIPLE PROPORTIONS: WHEN TWO ELEMENTS, A AND B, FORM MORE THAN ONE COMPOUND BY COMBINING WITH EACH OTHER, THE WEIGHTS OF ELEMENT B WHICH COMBINE WITH A GIVEN WEIGHT OF ELEMENT A STAND IN A RATIO WHICH CAN BE EXPRESSED BY SMALL WHOLE NUMBERS. It is especially worthy of note that the law was first stated before formulas were known and, indeed, was a major factor in first constructing formulas.

At about the same time (1809) that Dalton formulated his atomic theory, a French chemist, J. L. Gay-Lussac (1778-1850), observed a striking simplicity in the reactions of gases. He noted that, under conditions of constant temperature and pressure, the volumes of reacting gases, and their gaseous products, were in the ratio of small whole numbers.

A few years later, Amedeo Avogadro (1776-1856), an Italian physicist, explained Gay-Lussac's observation on the basis of Dalton's atomic theory. Avogadro's hypothesis was that equal volumes of gases, under the same conditions of temperature and pressure, contained the same number of molecules, (Section 2.6 & 2.7). Since the atoms in these molecules always remained whole when a reaction occurred, their rearrangement would be represented by simple whole number ratios. Since an equal number of molecules always occupied an equal volume, reacting volumes could also be expressed in small whole number ratios. Both Avogadro and Gay-Lussac made significant contributions to our knowledge of the structure of matter.

The atomic theory and the law of multiple proportions, as postulated by Dalton, were tested and accepted as correct, and for decades all scientific work was based on the assumption that atoms were indivisible.

Today there are major exceptions to some of Dalton's postulates. The most important exception is that the atom is divisible. (Dalton was still basically right in that atoms are not divided in ordinary chemical changes). The expression "splitting the atom" is a common one, and newspaper stories, television programs, and scientific publications have made us all aware that atoms are composed of three fundamental particles; electrons, protons, and neutrons.


The electrical nature of matter was observed over 2000 years ago by the Greeks. They discovered that objects could be electrified by rubbing. When two pieces of hard rubber rod are rubbed with fur, they acquire an electric charge and repel each other. When two pieces of glass rod are rubbed with a piece of silk, they also acquire an electrical charge and repel each other. However, when one of the charged hard rubber rods is brought near one of the charged glass rods, the two rods are attracted to each other. Since the rods exert a force on each other when separated, it is apparent that the region surrounding the charged rods is changed in some way. This region of influence that exerts a force on any charged object within the region is called an ELECTRIC FIELD. An electric field has no mass, can exist in a vacuum, and is always found in the space surrounding a charged body.

The above experiments provide evidence that there are two types of electric charge: negative and positive. They also illustrate the fundamental relationships governing charged bodies: like charges repel each other; unlike charges attract each other. The rubber rod is arbitrarily assigned a negative charge. That acquired by the glass rod is said to be positive. The force operating between the charged rods is called an electrostatic or coulombic force. The magnitude of the force between two charged objects was experimentally derived by Charles Coulomb (1736-1806), a French scientist.

The movement of charge through a conductor is known as an ELECTRIC CURRENT. In 1834, Michael Faraday (1791-1867) conducted ELECTROLYSIS experiments in which he used an electric current to decompose chemical compounds. For example, passing a current through a solution of silver nitrate (AgNO3) produces metallic silver. Faraday observed that the mass of an element obtained by the electrolysis of its compound was proportional to the quantity of electricity used and the atomic mass of the element. This observation suggests that some relationship exists between electricity and atoms.

In the late 19th and 20th centuries, a number of scientists devised experiments to study the nature of electric particles. These experiments consisted of studying the passage of electricity through gases in tubes known as GAS DISCHARGE or CROOKES TUBES. In 1879, Sir William Crookes (1832-1919) devised an apparatus consisting of a glass tube containing two metal plates, called ELECTRODES, which are connected to a source of electricity. The plate connected to the negative terminal of the source is called the CATHODE. The plate connected to the positive terminal is called the ANODE and consists of a movable metal cross which may be placed in either a horizontal or a vertical position. Most of the gas particles have been removed from the tube ( the tube has been partially evacuated) so that they will not interfere with the passage of electricity.

When the electric current is flowing and the anode is in a horizontal position, the glass opposite the cathode glows with a greenish light. This phenomenon is known as fluorescence. When the metal cross is placed in a vertical position, the rays traveling in a straight line from the cathode strike the cross but do not penetrate it. As a result, the shadow of the cross appears on the glass. This experiment shows that some kind of ray coming from the negative terminal travels in a straight line, transmits energy, and does not penetrate a metal sheet. Such a ray is called a CATHODE RAY.

To further determine the nature of cathode rays, a tube similar may be used. This tube has a metal screen which is coated with zinc sulfide and placed parallel to the path of the rays. Zinc sulfide is one of a number of substances which fluoresce when bombarded by cathode rays, thus permitting the detection of these rays. When the electrodes are connected to a source of high voltage electricity, the path of the rays is seen as a green line on the zinc sulfide screen. A magnet placed outside the tube causes the beam to be deflected downward. When the poles of the magnet are reversed, the beam is deflected upward. Interpretation of the effect of the magnetic lines of force (called a MAGNETIC FIELD) indicates that the cathode rays are composed of negatively charged particles. These particles were named ELECTRONS by G. J. Stoney (1826-1911), an English physicist, in 1874.

In 1897, J.J. Thomson (1856-1940), a British physicist, deflected a beam of cathode rays by applying an electric field of known strength. He then varied the magnetic field which was perpendicular to the electric field until the beam returned to its original undeflected position. Using the data from his experiment, Thomson was able to determine the ratio of the charge to mass e/m, of the electron and thus show that cathode rays consist of discrete particles of matter. By using different voltages, various metals for electrodes, and various gases in the tube, Thomson was able to show that the cathode rays always had the same properties. He concluded that ATOMS OF ALL SUBSTANCES CONTAIN THE SAME KIND OF NEGATIVE PARTICLES. In 1906, he was awarded the Nobel Prize.

In 1909, Robert A. Millikan (1868-1953), an American physicist, accurately determined the charge of the electron by his famous "oil drop" experiment. He measured the time required for electrically charged oil droplets to rise a measured distance when subjected to an electric field. He also measured the time required for the droplets to fall a specified distance in the absence of an electric field. He repeated these measurements with droplets containing various amounts of charge. Calculation of the charges on the droplets showed that they were always a multiple of the same small number. Since the values for the different charges on the droplets were all multiples of a fundamental charge, Millikan concluded that the fundamental charge was the absolute charge on a single electron. Knowledge of the charge determined by Millikan and the e/m ratio determined by Thomson enabled scientists to calculate the mass of the electron. The mass is approximately 9.1 X 10-28 grams or 1/1837 the mass of the hydrogen atom, the lightest of all atoms. For our practical purposes, the mass of the electron is negligible when compared with the total mass of an atom. For this work, Millikan was awarded the Nobel Prize in 1932.

The charge on an electron is designated by the symbol e-, and the superscript indicates that it bears a unit negative charge (-1). The charges on all other atomic particles are based on the -1 charge of the electron.


In 1886, Eugen Goldstein (1850-1930), used a perforated metal disc as the cathode of a tube. To his astonishment, he observed rays traveling in the direction opposite to the direction traveled by the cathode rays. These rays emerged from the holes on the side of the cathode away from the anode, and traveled in straight lines toward the end of the tube. Goldstein named them canal rays. Later, a German physicist, Wilhelm Wien, showed these rays to possess a positive charge, and calculated a rough value of the ratio of charge to mass for the particles composing the rays. Again, however, it was J.J. Thomson who really discovered the nature of the rays. He showed that the rays did consist of particles. He found that the least massive positive particles were emitted when hydrogen gas was used in the tube. When an electron is removed from a hydrogen atom, a single, positively charged particle remains. This particle is called a PROTON. It has a mass of 1.6725 X 10-24 grams. This is 1837 times the mass of an electron. This meant that the mass of a proton was almost the same as the mass of a hydrogen atom. The proton has become the standard unit of positive charge (+1). In 1914, H.G.J. Moseley (1887-1915), after completing his study of the X-ray beams emitted by different elements, stated "The ATOMIC NUMBER of an element is identified with the number of positive units of electricity contained in the atomic nucleus". The ATOMIC NUMBER of an element is equal to the number of protons in the nucleus. This is also equal to the number of electrons outside the nucleus of each uncharged atom of the element.


The first conclusive evidence that atoms are complex rather than "indivisible," as stated in the atomic theory, of John Dalton, came with the discovery of radioactivity by Antoine-Henri Becquerel 1852-1908), a French physicist, in 1896. The term RADIOACTIVITY applies to the spontaneous decomposition of atoms of certain elements, such as radium and uranium, in-to simpler elements, and the simultaneous production of one or more of three kinds of radiations.

These three kinds of rays may be studied by placing samples of the radioactive material in the bottom of a narrow hole bored in a block of lead, and allowing the emitted rays to pass through strong electric fields. One type of ray is curved toward the negative part of the electric field and so must consist of positively charged particles. These are called ALPHA PARTICLES. They are not protons as shown by the fact that they have a mass of 4 and a charge of +2. Experiment has shown that they are helium ions, helium atoms that have lost two electrons. A second type of ray is deflected toward the positive part of the field, and the path is bent much more than that of the alpha particles. These facts indicate that these particles are negatively charged and much smaller in mass than the alpha particles. They are called BETA PARTICLES. Beta particles are electrons with velocities approaching that of the speed of light (about 186,000 miles/second). They are believed produced from the nucleus rather than from the electron cloud. The third kind of ray is not deflected when it passes through an electric field. These rays are more penetrating than alpha or beta particles and are called GAMMA RAYS. Gamma rays are electromagnetic waves of very short wavelength. (See Chapter 5)


A third particle remained undetected for a long time. Its existence had been predicted by Ernest Rutherford (1871-1937), (Lord Rutherford), an English physicist, in 1920, but the first evidence of the particle was not observed until 1930 by Walter Bothe (1891-1957). Bothe was conducting some experiments in radioactivity, and one of his tests failed to produce the expected results. He expected the material with which he was working to emit protons. Instead, he obtained high energy rays which he thought were gamma rays.

Another English scientist, James Chadwick (1891-1974), duplicated bothe's work in 1932 and found that the material was not emitting gamma rays but high energy particles with no charge and with essentially the same mass as the proton. Through his work, Chadwick is credited with the discovery of the NEUTRON. The mass of the neutron is now known to be 1.6748 x 10-24 grams. The neutron plays the leading role in the drama of nuclear energy. It serves as the key which unlocks the energy of the nucleus and produces the radioisotopes which are used in medicine, industry, agriculture, and research.

Scientists do not believe that a nucleus contains individual protons and neutrons. Rather, they picture the nucleus as being composed of neutrons and protons which lose their identity as they merge to form a tiny, dense ball. The total charge on the nucleus is equal to the number of protons, and its total mass is approximately equal to the sum of the masses of its neutrons and protons.



As noted earlier, protons and neutrons each contribute very nearly one unit of atomic mass to atoms. This being so, we would expect to find the atomic weights (mass) of the elements to be approximately whole numbers. The fact that many atomic weights are far from being whole numbers led to the discovery that most elements exist as mixtures of two or more kinds of atoms of different atomic masses, but of similar chemical properties.

While working with neon, J.J. Thomson observed what appeared to be two kinds of neon atoms. They were exactly alike chemically, but differed slightly in mass. This same phenomenon had been observed earlier by Frederick Soddy (1877-1956), an English chemist, while working with radioactive materials. Soddy coined the word ISOTOPE for different kinds of atoms of the same element and Thomson adopted Soddy's term. Today it is known that there are isotopes of every element. ATOMS WHICH POSSESS THE SAME NUMBER OF PROTONS BUT CONTAIN A DIFFERENT NUMBER OF NEUTRONS ARE CALLED ISOTOPES.

Chlorine, with an atomic mass of 35.453, exists as two kinds of chlorine atoms. These have masses very close to the whole numbers 35 and 37. Both types of chlorine atoms have an atomic number of 17, which means that they each have 17 protons in the nucleus. The difference must lie, therefore, in the number of neutrons in the nuclei of the different types of atoms; chlorine-35 has 18 neutrons and chlorine-37 has 20 neutrons. It is important to remember that the only difference in composition between isotopes of the same element is in the number of neutrons in the nucleus. The atomic mass of an element is an average of the weights of the isotopes of the element in the proportions in which they normally occur in nature. The atomic weight of 35.453 for chlorine indicates that the atoms of weight 35 are more abundant than the atoms of weight 37.

In Section 2.5, it was noted that a carbon isotope is the present standard in the atomic weight scale. Carbon has three principal isotopes, carbon-12, carbon-13, and carbon-14. Each has six protons in the nucleus and six electrons in the electron cloud. Carbon-12 has in addition, six neutrons in the nucleus, carbon-13 has seven neutrons, and carbon-14 has eight neutrons. The atomic weight scale is based up-on the arbitrarily assigned value of the exact number 12 for the mass of the lightest and most abundant carbon isotope, Carbon-12. The three isotopes of carbon occur in such proportion in nature that the average weight of a mixture is 12.01115. Thus, the weights as represented on the PERIODIC TABLE OF ELEMENTS for each element is an average mass of all the isotopes of that element as it occurs in nature.

John Dalton had assumed that atoms were indivisible. The discovery of subatomic particles led to a major revision of his atomic theory. The discovery of isotopes led to a second major revision of Dalton's atomic theory, in which all atoms of an element were exactly alike (Section 3.4). Dalton's theory was not discarded, it was merely modified to include the additional information.


During the period, 1912-1913, Lord Rutherford had assembled in his laboratory a remarkable team of physicists, including Niels Bohr (1885-1962), a young Dane. Experiments were performed under Rutherford's direction which succeeded in showing that the atom consisted of a central, positively charged nucleus, surrounded in some manner by electrons. Since the electrons were negatively charged and attracted to the larger positive, it seemed that there was nothing to prevent the electrons from falling into the nucleus. The discussions among these physicists, led by Rutherford and Bohr, resulted in the proposal that the electrons are in "orbit" around the nucleus, in much the same manner as the earth is in orbit around the sun. They postulated that there is a force of attraction between the electrons and the nucleus, just as there is a force of attraction between the earth and sun. The Rutherford-Bohr model of the atom attributed the same situation to electrons in orbit around the nucleus, as exists for the earth in orbit around the sun. Such a picture of the atoms is sometimes called the planetary atomic model.


For our purposes, the results of the Rutherford-Bohr theory indicates that the hydrogen atom consists of one proton (as a nucleus) and one electron moving about the proton. Larger atoms have several electrons moving about a positive nucleus which consists of several protons and neutrons. However, the exact path of the electron in its particular energy level cannot be determined. The German physicist Werner Karl Heisenberg (1901-1976) expressed this in a form which has come to be known as the HEISENBERG UNCERTAINTY PRINCIPLE: IT IS IMPOSSIBLE TO DETERMINE ACCURATELY BOTH THE MOMENTUM AND THE POSITION OF AN ELECTRON SIMULTANEOUSLY. (The momentum is the mass of the particle multiplied by its velocity.)

The electrons are located in definite energy levels. Inasmuch as the physical and chemical properties of an atom are primarily determined by the energies associated with the electrons in their movement within the atom, the energy of the electron is much more important to the chemist than the actual position of the electron. Some chemists prefer to consider the electron in terms of a cloud of negative charge (electron cloud), with the cloud being dense in regions of high probability of electrons and more diffuse in regions of low probability. The ELECTRON CLOUDS are spoken of as ATOMIC ORBITALS.


The nuclei of atoms which are more complex than that of hydrogen are surrounded by orbital electrons arranged in a series of SHELLS, or ENERGY LEVELS. The closer an energy level is to the nucleus, the more strongly the electrons in that level are attracted to the nucleus because the electrostatic force of attraction between particles of opposite electrical charge increases rapidly with decreased distance of separation. The orbital electrons are classified in terms of their energy levels, i.e., on the basis of the firmness with which they are abound to the nucleus in the atom. The energy levels or shells are designated by the numbers 1, 2, 3, 4, 5, 6, and 7 or by the letters K, L, M, N, O, P, and Q, starting with the one nearest the nucleus. The maximum number of electrons found in any given energy level or shell is as follows: K=2, L=8, M=18, N=32, O=32, P=18, and Q=8.

Spectroscopic data reveal that the electrons in the principal energy levels are distributed in energy sublevels. The energy sub-levels are designated in order of increasing energy as s, p, d, and f. The maximum number of electrons found in any given sublevel is as follows: s=2, p=6, d=10, and f=14. The first principal energy level (K) contains only an s sublevel; the second (L) contains s and p sublevels; the third (M) contains s, p, and d sublevels; and the fourth (N) contains s, p, d, and f sublevels. Energy sublevels are also referred to as energy subshells.

These subshells are further divisible into ORBITALS, each of which can contain a maximum of two electrons. Thus, an s subshell, which is made up of one orbital, can contain a maximum of two electrons; a p subshell has three orbitals and can contain up to six electrons; a d subshell has five orbitals and can contain a maximum of ten electrons; and an f subshell has seven orbitals and can contain a maximum of fourteen electrons.

It had been found by Wolfgang Pauli (1900-1958) that all of the electrons spin either clockwise or counterclockwise, and that the two electrons of a given orbital are identical in all respects except that they are spinning in opposite directions. According to the PAULI EXCLUSION PRINCIPLE, NO TWO ELECTRONS IN THE SAME ATOM MAY HAVE IDENTICAL ENERGIES. Two electrons of an atomic orbital having opposite spins prevents them from having identical energies.

1. Electrons enter the lowest sub-orbital of the lowest energy shell first.

2. No electron pairing takes place in the p, d or f sub-orbitals until each sub-orbital of the given set contains one electron.

This is known as HUND'S RULE. For example, each of the three p sub-orbitals in the second major energy level (L) receives one electron before any of them receives a second electron. This implies that it is more difficult for an electron to enter a sub-orbital that already contains an electron than to enter an unoccupied sub-orbital.

3. No sub-orbital can have more than two electrons.

K L M N O P Q energy levels
1s 2s2p 3s3p3d 4s4p4d4f 5s5p5d5f 6s6p6d 7s7p Subshells
2 2 6 2 6 10 2 6 10 14 2 6 10 14 2 6 10 2 6  
2 8 18 32 32 18 8 Total
Hydrogen H 1s1 1 electron
Lithium Li 1s2/2s1 3 electrons
Chlorine Cl 1s2/2s22p6/3s23p5 17 electrons
Bromine Br 1s2/2s22p6/3s23p63d10/4s24p5 35 electrons

As noted, when determining the electronic structure of atoms of the various subshells, it is convenient to consider the subshells which electrons would enter if these atoms were built up in order of increasing atomic number, beginning with hydrogen. As each additional electron enters the atom, it will tend to occupy the available orbital of lowest energy, and electrons enter higher energy orbitals only after lower energy orbitals have been filled to capacity.


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Copyright May 1987 James R. Fromm (jfromm@3rd1000.com)   (Revised March 2000)