Early Ideas of Atomic Structure
As mentioned in the previous section titled “A Short History of Chemistry,” many scien- tists identified elements, determined their characteristics, similarities, and differences, and designed symbols for them. Using unique experiments, scientists devised ways to define the structure of atoms and determine atomic weights, sizes, and electrical charges as well as energy levels for atoms.
Many of these men and women recognized the existence of some order in the manner in which chemicals relate and react to each other. Although these scientists could not see the atoms themselves, they were aware that the structure of each element’s atoms has something to do with these characteristics. There were several attempts to organize the elements into a chart that reflected the particular nature of the atoms for these elements. Before the periodic table of the chemical elements was developed as we know it today, several relationships had to be established. (See the next section for more on the periodic table of the chemical elements.)
The concept of electrons had been known for many years, but determining how these negatively charged particles react required experimentation and analysis of data. In about 1897, Joseph John Thomson (1856–1940) sent streams of electrons through magnetic fields, which resulted in the dispersion or spreading of the electrons. Thomson’s experiments, and those of others, led him to speculate that the atom was a positively charged “core” and that negatively charged particles of energy surrounded and matched the positive charge of this core or nucleus. Further, when these electrons were excited or “stirred up” with strong light, electricity, or mag- netism, some of them were driven from the outer regions of the atom. This was one of the first experimental evidences for the structure of the atom. Many refinements of this concept were made by the work of several scientists, including the French chemist Marie Sklodwska Curie (1867–1943), the British physicist Baron Ernest Rutherford (1871–1937), the American physi- cist Robert Andrew Millikan (1868–1953), and others of many nationalities.
Two of Rutherford’s students conducted a classic experiment to determine the structure of an atom. They beamed alpha particles through a sheet of gold foil that was 1/50,000 of an inch thick. Since the thickness of this thin foil was only about two thousand atoms of gold, it
should have allowed all the alpha particles to pass through and be detectable on the other side of the foil. The experimenters recorded the pattern of alpha particles with a detecting instru- ment located behind the foil. Rutherford noticed that most particles went straight through the foil as if nothing interfered with them. However, unexpectedly, a few seemed to be diverted from the target. As it turned out, one out of every 10,000 alpha particles, as it passed through the foil, was deflected sideways, away from the center of the target located behind the foil, similar to a billiard ball glancing off another ball when struck. During the experiment, one out of every 20,000 particles bounced back from the foil toward the source of the alpha particles and thus did not pass through the gold foil at all. This indicated that some alpha particles were being deflected by something in the foil. After making some calculations, Rutherford concluded that this backward and side scattering of the alpha particles was evidence of a few collisions with something that had almost all its mass concentrated in a central, very small “nucleus.” He also determined that this tiny nucleus had a positive charge. Since the vast majority of the alpha particles passed straight through the foil to hit the center of the target used to detect the particles, the atoms of gold must be composed of mostly empty space. He stated, “It was almost as incredible as if you fired a 15-inch shell at a piece of tissue paper, and it came back to hit you.”* One way to understand the great distance between an atom’s
nucleus and its surrounding electrons is to imagine a baseball on the floor of your living room as the nucleus of an atom. The electrons that surround the nucleus are about the size of peas that are orbiting in three dimensions at about a 10-mile radius from the baseball. This means that an atom with a 20-mile diameter would have a nucleus about the size of a baseball. The conclusion may be that there is more empty space than there is solid matter in matter and the universe, including in you.
Robert Andrew Millikan (1868–1953) continued the work of J. J. Thomson by conducting a classic experiment to determine the exact charge and mass of the electron. He constructed two flat round brass plates with a one-centimeter space between them. He drilled a tiny pin- hole through the top plate and arranged them so that the top plate carried a positive charge and the bottom plate a negative charge. Placing a microscope at the edge of the opening between the plates, he then atomized thin oil droplets over the top plate. As the oil droplets drifted through the tiny hole and moved between the plates, Millikan measured the time it took the droplets to descend from the top to the bottom plate. By varying the electric charge on the droplets, he could control their rate of descent across the space between the charged plates. This enabled him to determine the negative electrical charge of electrons. His figure was close to the accepted figure of –1.602177 × 10-19 coulombs. From this data, he then cal-
culated the mass of the electron to be 1/1,835 that of the hydrogen atom. The hydrogen atom is essentially composed of a single proton with a single negative charge, or electron.
Some questions still required answers before a logical organizational chart of the elements could be completed, among them how to determine the energy, position, and number of elec- trons in atoms of different elements. When electrons are “excited” by an input of energy, they jump to a higher energy level, orbit, or shell—that is, from a shell (energy level) that is closer to the nucleus to a higher energy level further from the positive nucleus—and then they return to their original energy level closer to the nucleus, or they are driven off the outer region of * Rutherford, Ernest. Qtd. in Aaron J. Ihde. The Development of Modern Chemistry. New York: Dover Publications, 1984.
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the atom. When the energized electrons return from their higher energy position to their original lower energy position, they emit the acquired energy in unique electromagnetic frequencies, across the elec- tromagnetic spectrum. This energy release occurs in uniform “packets” of energy known as photons (quantum units of electromagnetic radiation). This jump of the electron(s) changes the atom and is known as a quantum transition or a quantum leap of radiant energy, units of which cannot be divided into smaller bits. Contrary to the media’s use of “quantum leap” when referring to some large change
or action, a quantum leap is a very, very small bit of matter or energy.
Max Karl Ernst Ludwig Planck (1858–1947), following some theories of electromagnetic radiation, devised a mathematical equation to explain the relationship between energy and wavelengths of light. Now referred to as Planck’s constant, it is an invariant that, when multi- plied by the frequency of electromagnetic radiation, describes the amount of energy contained in one quantum. It is an extremely small measurement, but the concept is so fundamental that it heralded the beginning of modern physics in the early 1900s.
Werner Karl Heisenberg (1901–1976), the founder of quantum mechanics, described the behavior of very small particles and their interactions based on some higher mathematical con- cepts. The theory is somewhat easier to understand in a descriptive sense. His most unusual theory changed classical physics and chemistry and is now referred to as indeterminacy or the “uncertainty principle.” In essence, it states that it is impossible to know both the exact posi- tion and the momentum (mass) of a subatomic particle at the same moment. Thus, the ulti- mate nature of matter is not susceptible to objective measurement (indeterminacy). However, an accurate prediction (statistical average) of the nature of matter is possible when based on statistical probabilities. An interesting aspect of the Heisenberg uncertainty principle is that merely viewing and attempting to measure subatomic particles may alter their positions and other characteristics. It might be mentioned that this physics concept of quantum uncertainty has nothing to do with our macro-world’s philosophical concepts of being “uncertain” or of sociological relativity.
Wolfgang Pauli (1900–1958), an American physicist, was awarded a Nobel Prize in 1945 for developing the “exclusion principle.” In essence, it states that a particular electron in an atom has only one of four energy states and that all other electrons are excluded from this electron’s energy level or orbital. In other words, no two electrons may occupy the same state of energy (or position in an orbit around the nucleus). This led to the concept that only a cer- tain number of electrons can occupy the same shell or orbit. In addition, the wave properties of electrons are measured in quantum amounts and are related to the physical and, thus, the chemical properties of atoms. These concepts enable scientists to precisely define important physical properties of the atoms of different elements and to more accurately place elements in the periodic table.
The terms “shell,” “orbital,” and “energy level” are sometimes used interchangeably. In this text we use the term “shells” most often because it is descriptive and conveys the image of three-dimensional layers or structures of electrons surrounding the nucleus. The term
Figure 2.1: Rutherford’s experimental apparatus to determine the structure of the atom.
“orbit” is more closely related to the image of two-dimensional concentric rings, similar to an archery target. The term “energy level” describes the energy that electrons possess, depending on their distance from the atom’s nucleus. The term “orbital” is the more descriptive term and describes the original four distinct energy levels detailed in Pauli’s exclusion principle based on the angular momentum quantum numbers. Because the original four energy levels for electrons surrounding the nucleus were identified by Wolfgang Pauli and others, three more orbitals (energy levels) have been added to accommodate new artificial elements as they have been discovered. Shells, orbitals, and energy levels make up the levels of electron energy sur- rounding the nuclei. Similar to “shell,” “orbital,” and “energy level,” the notions of subshells or orbitals are used interchangeably. This electron structure is depicted by the figure for each element listed in the section, “Guide to the Elements.” The patterns for electron configuration for each of the following seven energy levels (shells or orbits) follow:
1. The first shell or energy level out from the nucleus is called the “K” shell or energy level and contains a maximum of two electrons in the “s” orbital—that is, K = s2, where the “K” represents the shell number (or principle quantum number), the “s” describes the orbital shape of the angular momentum quantum number, and the “2” is the maximum number of electrons that the “s” orbital can contain. This particular sequence is “K = s2,” which means K shell contains 2 electrons in the “s” orbital. This is the sequence for the element helium. Look up helium in the text for more information.
2. The second shell or energy level is “L” and may contain a maximum of eight electrons; its orbital is called “p” and can contain a maximum of six electrons. Therefore, its sequence would be K = 2 (s2) and L = 8 (s2 + p6). This sequence of 10 electrons in the first two shells (K and L) represents the element neon, which has an atomic number of 10. 3. The third shell or energy level is “M” and may contain a maximum of 18 electrons; its
orbital is called the “d” subshell, and it may have a maximum of 10 electrons: for example, K = 2 (s2), L = 8 (s2 + p6), and M = 18 (s2 + p6 + d10). See the elements in the text for more details.
4. The fourth shell or energy level is “N,” which may contain a maximum of 32 electrons; its orbital is called “f” subshell: for example, K = 2 (s2), L = 8 (s2, p6), M = 18 (s2, p6, d10), and N = 32 (s2 + p6 + d10 + f14). Refer to the elements in the text for more information. 5. The fifth shell or energy level is “O,” also with a possible maximum of 32 electrons; its
highest orbital is also the “f” subshell: K = 2 (s2), L = 8 (s2, p6), M = 18 (s2, p6, d10, f14), and O = 32 (s2, p6, d10, f14).
6. The sixth shell or energy level is “P,” with a maximum of 10 electrons. Remember that the numbers of electrons in each shell (K, L, M, N, O, and P) for each individual element are added together to find the total number for a particular atom, which is also the element’s atomic number.
7. The seventh shell or energy level is “Q,” with a maximum of two electrons, and is repre- sented as Q = 2 (s2) with just two electrons in it first orbital. (This sequence holds until the element ununtrium-113, where Q = 3 (s2, p1), and those heavy elements beyond 113, where the “Q” shell may contain more than three electrons.)
Note: All elements with a depiction of higher energy levels or shells are both syntheti- cally produced and radioactive, and electrons are added to inner shells rather than the usual
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outer shell of the elements. As these heavy elements’ atomic numbers increase, their half-lives decrease to a fraction of a second, and they are produced one or few atoms at a time. This concept is depicted with the data presented for each element.
Some Theoretical Atomic Models
The earliest concept of atomic structure dates back to Greece in the fifth century BCE, when Leucippus and Democritus postulated that tiny particles of matter, which they called atomos,
were indivisible.
Over the centuries, many other concepts were proposed to explain the nature of matter— many of them extensions of the Greek concept of an ultimately indivisible and indestructible elementary bit of matter. But it was not until J. J. Thomson proposed his model of the atom, consisting of a sphere with an agglomeration of particles with negative electric charges some- how positioned randomly inside a very small ball of matter, that the modern structure of the atom began to take shape.
The Rutherford model of the atom is a significant improvement over the Thomson model. Baron Ernest Rutherford incorporated the background and understanding of many scientists as he developed experiments designed to show that the atom has a central and small, but relatively heavy, nucleus. His experiments verified that this positively charged dense nucleus has negatively charged electrons surrounding it at a very great distance compared to the size of the nucleus. This concept resembles the planets revolving around the sun, including with regard to the laws of motion and energy.
The Bohr model of the atom took shape in 1913. Niels Bohr (1885–1962), a Danish physicist, started with the classic Rutherford model and applied a new theory of quantum mechanics to develop a new model that is still in use, but with many enhancements. His assumptions are based on several aspects of quantum theory. One assumption is that light is emitted in tiny bunches (packets) of energy call photons (quanta of light energy).
The Bohr model continues:
First, the orbital or quantum theory of matter assumes that the electron is not a particle, as we normally think of particles. Orbital theory considers the electron as a three-dimensional wave that can exist at several energy levels (orbitals), but not at the same time.
Second, the electrons are in constant motion around the nucleus, even though it is not easy to determine the position of a particular electron in its shell at any particular moment.
Third, the electrons are revolving at different distances from their nuclei; these pathways are called shells, orbitals, or energy levels.
Fourth, there are more than a few shells, orbitals, or energy levels in most atoms.
Fifth, the electrons can continue to move in a specific shell without emitting or absorbing energy.
Sixth, an electron at a specific energy level will remain in its shell until it either loses energy and “jumps” to an outer shell or gains energy and proceeds to a higher energy level (shell). Seventh, if the electron is excited by external energy, it can “jump” to a different shell or higher energy level.
Eighth, when the electron returns to its former shell or lower energy level, it will emit the energy (photon), which represents energy the electron acquired to raise it to the higher energy level.
Bohr’s idea led to the comparison that likened the structure of the atom to the structure of an onion. The outer layers of skin on an onion are the “shells” where the electrons exist.
At the center of the onion is a dense, tiny BB-like shot for a nucleus. Unlike the onion, however, the area between the nucleus and the electron’s shells is merely space. There is no distinct boundary for the shells. The electrons assigned to a particular shell are in constant motion, so the shell does not seem to have a sharp definition. It is something like a “fuzzy ball,” with no distinct edges.
A similar model proposed at that time was related to “raisins in a plum pudding,” with the electrons resembling raisins spread throughout the pudding. The fuzzy-ball or raisin-in-plum-pudding model of the atom has no sharp boundaries, but does have definite limits as to the number and energy levels for the electrons residing within each shell. The fuzzy ball depicts the energy concept of the “electron cloud,” which considers the electrons as energy levels around the nucleus. The concept states that the atom is spherical, and the electrons are all over the atom at any one time. The electron cloud concept statistically depicts the probable distribution of electrons as they exist at any particular distance from the nucleus at any one particular time. The fuzziness is the image one might see while viewing the atom over an extended period, such as a time exposure with a camera.
The number of electrons in each shell is partially dependent on how far the shell is from the nucleus. In addition, electrons assigned to a specific shell stay in position until forced out by an input or loss of energy. Even more important, a particular shell or energy level “dis- likes” having any electrons missing. If a shell does not have its complete quota of electrons, the atom will “bond” with other atoms by “taking in electrons” or “giving up electrons” or