SAT Subject Test Chemistry




Atomic Structure and the Periodic Table of the Elements


Order of Filling and Notation

The sublevels do not fill up in numerical order, and the pattern of filling is shown on the right side of the approximate relative energy levels chart (Figure 9). In the first instance of failure to follow numerical order, the 4s fills before the 3d. (Study Figure 9 carefully before going on.)

Figure 9. Approximate Relative Energy Level of Subshells

Figure 9. Approximate Relative Energy Level of Subshells


1s2 2s2 2p6 3s2 3p6 4s1


1s2 2s2 2p6 3s2 3p6 4s2


1s2 2s2 2p6 3s2 3p6 4s2 3d1 (note 4s filled before 3d)

There is a more stable configuration to a half-filled or filled sublevel, so at atomic number 24 the 3d sublevel becomes half-filled by taking a 4s electron;


1s2 2s2 2p6 3s2 3p6 3d5 4s1

and at atomic number 29 the 3d becomes filled by taking a 4s electron:


1s2 2s2 2p6 3s2 3p6 3d10 4s1

Table 3 shows the electron configurations of the elements. A triangular mark  indicates an outer-level electron dropping back to a lower unfilled orbital. These phenomena are exceptions to the Aufbau Principle. By following the atomic numbers throughout this chart, you will get the same order of filling as shown in Figure 9.

Table 3. Electron Configuration of the Elements

Note: Follow order of the atomic numbers to ascertain the order of filling.

By following the atomic numbers in numerical order in Table 3 you can plot the order of filling of the orbitals for every element shown.

A simplified method of showing the order in which the orbitals are filled is to use the following diagram. It works for all the naturally occurring elements through lanthanum, atomic number 88.


This is a simple way to remember the order of filling orbitals.

Start by drawing the diagonal arrows through the diagram as shown. The order of filling can be charted by following each arrow from tail to head and then to the tail of the next one. In this way you get the same order of filling as is shown in Figure 9 and Table 3:

1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s2 4f14 5d10 6p6 7s2

Lewis Structures (Electron Dot Notation)

In 1916 G. N. Lewis devised the electron dot notation, which may be used in place of the electron configuration notation. The electron dot notation shows only the chemical symbol surrounded by dots to represent the electrons in the incomplete outer level. Examples are:

The symbol denotes the nucleus and all electrons except the valence electrons. The dots are arranged at the four sides of the symbol and are paired when appropriate. In the examples above, the depicted electrons are the valence electrons found in the outer energy level orbitals.

4s1 is shown for potassium (K)

4s2 4p3 are shown for arsenic (As)

5s2 is shown for strontium (Sr)

5s2 5p5 are shown for iodine (I)

6s2 6p6 are shown for radon (Rn)

Noble Gas Notation

Another method of simplifying the electron distribution to the orbitals is called the noble gas notation. In this method you represent all of the lower filled orbitals up to the closest noble gas. By enclosing its symbol in brackets, it represents all of the complete noble gas configuration. Then the remaining orbitals are written in the usual way. An example of this can be shown by using the third period of elements. By using neon as the noble gas, you write [Ne] to represent its orbital structure, which is 1s2 2s2 2p6. This allows you to write an element like sodium as [Ne] 3s1, which is called sodium’s noble gas notation. The table in the next section shows the noble gas notations of some of the transition elements in the fourth period of elements. Notice that the base structure of argon is used and represented as [Ar].