JACK W. EICHINGER, Jr. The Florida State University, Tallahassee, Florida
T I M E can be saved and confusion avoided by developing a systematic chart of the elements based on the energy levels of atomic orbitals very early in the general college chemistry course. Excellent results have been obtained by substituting this procedure for the usual historical development of the periodic table. The resulting chart is used in place of a periodic table throughout the year and resembles the usual tables sufficiently so as to present no problems when they are encountered later. A description of the chart has been puhli~hed.~This paper will describe how students may prepare their own charts and gain some understanding of electron configurations and families of elements. The theoretical basis for this type of chart has been adequately presented by Longuet-Higgiw3 Beginning students, however, are spared these details in favor of the simplified approach of Timm.& As a first approximation, atomic orbitals are visualized as "locations" which electrons may occupy. No two orbitals have exactly the same energy. The energy referred to in this case is the potential energy of an electron occupying the orbital. By adding energy an electron may be promoted to a higher energy level or "location" which we will visualize initially as being a little further removed from the positive nucleus. This concept is easily grasped by beginners whereas standing waves and dumbbell-shaped orbitals leave them hopelessly confused.
BAR MAGNETS DEMONSTRATE ELECTRON PAIRS
The fact that each orbital may be occupied by two electrons with opposed spins can be presented in a readily acceptable manner by the use of bar magnets. Electrons possess an electrical charge; moving charges generate a magnetic field; reversing the motion reverses the field. Bar magnets with their north and south poles painted in contrasting colors may be pushed about the lecture table and will readily form "electron pairs" when properly orientated. Energy is required to "uncouple" a pair of electrons. This may be demonstrated by twisting one of the magnets representing the magnetic fields about an electron pair, a procedure which obviously requires the application of force. After the magnet has been turned through a half-circle, the "uncoupled" pair will fly apart when released. A simple mnemonic will help the beginner to deter-
' Presented before the Division of Chemical Education at the 131st Meeting of the American Chemical Society, Miami, April, 1957. EICHINGER, J. W., J. CHEM.EDUC., 34, 7&1 (1957). S L ~ ~ ~H. C., ~ J. CHEM. ~ ~ EDUC., - 34, H 3Wl ~ (1957). ~ ~ 'TIMM,J. A,, "General Chemistry," 3rd ed., McGraw-Hill Book Co., Ine., New York, 1956, p. 84.
mine the composition of the electron shells and subshells-"sober physicists don't find giraffes hiding in kitchens." We call this the "Key to the Orbitals." The source is unknown and its only virtue is that it is silly enough to be easily remembered. All the student needs to do is count by ones and then by twos. Counting by ones gives the types of orhitals found in each main shell. Counting by twos gives the number of each type of orbital present in a particular shell. Counting by ones shows that the first shell contains only s orhitals; the second, s and p; the third. s, p, and d; the fourth, s, p, d, and f; etc. Starting again at one and counting by twos indicates that a shell may contain only one s orbital, three p orbitals, five d orhitals, seven f orbitals, etc. Using these simple rules the student can easily figure out that the eighth shell contains fifteen k orbitals. He will feel bett,er after you tell him that these higher levels are populated only during experiments involving extremely high energies and need not concern the chemist who is usually interested in the "ground states" of the atoms. CHART BASED ON RELATIVE ENERGIES OF SUBSHELLS
It is helpful to provide mimeographed blank charts for classroom use. These resemble the one displayed on the wall except .that all numbers and letters are deleted. The figure shows a partially filled form. Each orbital is represented by two boxes, one for each electron. Electron subshells are aligned vertically while the horizontal main electron shells slope downward to the right to indicate the overlapping of the energy levels. For example, the 3d orbitals lie at a higher energy level than the 4s and therefore do not begin to fill with electrons until after the 4s orbital is filled. Reference to the mnemonic enables the student to label these subshells properly. In developing the chart, one must imagine a positive nucleus located above and somewhat left of the center of the chart area. It exerts a strong attraction for the negatively charged electrons which in reality are distributed throughout three-dimensional space, but which are to be tabulated in the flat chart area.for convenience. This chart area, of course, bears no resemblance to the geometrical pattern or the scale of distances between electrons. It does suggest in a helpful manner the additional energy that must be expended to force an electron into a "position" farther away from the positive nucleus (downward on the chart). This can be visualized by stretching a rubber band from the imaginary location of the nucleus above the chart to the ~ "position" ~ ~ of , the electron. Chart "positions" do not imply that the electrons occupy fixed positions in space. They correspond instead to fixed energy levels. JOURNAL OF CHEMICAL EDUCATION
Hydrogen has one electron which in the ground state will occupy the orbital of lowest energy, Is. We place the atomic number and symbol for hydrogen in the upper left box. The two electrons of helium completely fill the 1s orbital so ,He is placed in the box to the right. We continue to build up the chart by assigning all of the electrons possessed by an atom to t.he lowest energy levels available. The law of maximum multiplicity is first encountered in the 2 p subshell. This formidable-sounding maxim merely states that no orbital may contain more than one electron until all the other orbitals in the particular subshell contain at least one. The chart provides for this automatically by separating the boxes into two groups by means of double vertical lines. Thus, in the 2 p subshell we place sBin the first box, GCin the second
VOLUME 34, NO. 10, OCTOBER, 1957
and TN in the third. The double line has now been reached and each of the three 2 p orbitals contains one electron. Additional electrons can be placed in this subshell only if they pair up with those already present. When we place in the first box to the right of the double line, its eighth electron has paired with another to completely fill one of the 2 p orbitals leaving single electrons in the other two. The 2 p subshell of 9Fcontains two pairs of electrons and a single unpaired electron in the third orbital. The value of this kind of information later in the course is readily apparent. A portion of the chart is built up during the class period and the students complete it at home. A4dditional information may be added to the chart from time to time but it is preferable to wait until an application presents itself.
