how to write electronic configurations

Introduction

There are four principle orbitals (s, p, d, and f) which are filled according to the energy level and valence electrons of the element. All four orbitals can hold different number of electrons. The s-orbital can hold 2 electrons, and the other three orbitals can hold up to 6, 10, and 14 electrons, respectively. The s-orbital primarily denotes group 1 or group 2 elements, the p-orbital denotes group 13, 14, 15, 16, 17, or 18 elements, and the f-orbital denotes the Lanthanides andActinides group. The main focus of this module however will be on the electron configuration of transition metals, which are found in the d-orbitals (d-block).

The electron configuration of transition metals is special in the sense that they can be found in numerous oxidation states. Although the elements can display many different oxidation states, they usually exhibit a common oxidation state depending on what makes that element most stable. For this module, we will work only with the first row of transition metals; however the other rows of transition metals generally follow the same patterns as the first row

The Electron Configurations of Atoms

The electron configuration of an atom shows the number of electrons in each sublevel in each energy level of the ground-state atom. To determine the electron configuration of a particular atom, start at the nucleus and add electrons one by one until the number of electrons equals the number of protons in the nucleus. Each added electron is assigned to the lowest-energy sublevel available. The first sublevel filled will be the 1s sublevel, then the 2s sublevel, the 2p sublevel, the 3s, 3p, 4s, 3d, and so on. This order is difficult to remember and often hard to determine from energy-level diagrams such as Figure 5.8
A more convenient way to remember the order is to use Figure 5.9. The principal energy levels are listed in columns, starting at the left with the 1s level. To use this figure, read along the diagonal lines in the direction of the arrow. The order is summarized under the diagram.

FIGURE 5.9 The arrow shows a second way of remembering the order in which sublevels fill.

An atom of hydrogen (atomic number 1) has one proton and one electron. The single electron is assigned to the 1s sublevel, the lowest-energy sublevel in the lowest-energy level. Therefore, the electron configuration of hydrogen is written:

For helium (atomic number 2), which has two electrons, the electron configuration is:

He: 1s²

Two electrons completely fill the first energy level. Because the helium nucleus is different from the hydrogen nucleus, neither of the helium electrons will have exactly the same energy as the single hydrogen electron, even though all are in the 1s sublevel. The element lithium (atomic number 3) has three electrons. In order to write its electron configuration, we must first determine (from Figure 5.9) that the 2s sublevel is next higher in energy after the 1s sublevel. Therefore, the electron configuration of lithium is:

Li: 1s²2s¹

Boron (atomic number 5) has five electrons. Four electrons fill both the 1s and 2s orbitals. The fifth electron is added to a 2p orbital, the sublevel next higher in energy (Figure 5.9). The electron configuration of boron is:

B: 1s²2s²2p¹

Table 5.2 shows the electron configurations of the elements with atomic numbers 1 through 18. The electron configurations of elements with higher atomic number can be written by following the orbital-filling chart in Figure 5.9.

TABLE 5.2 Electron configurations of the first 18 elements

Element	Atomic number	Electron configuration
hydrogen	1	1s1
helium	2	1s2
lithium	3	1s22s1
beryllium	4	1s22s2
boron	5	1s22s22p1
carbon	6	1s22s22p2
nitrogen	7	1s22s22p3
oxygen	8	1s22s22p4
fluorine	9	1s22s22p5
neon	10	1s22s22p6
sodium	11	1s22s22p63s1
magnesium	12	1s22s22p63s2
aluminum	13	1s22s22p63s23p1
silicon	14	1s22s22p63s23p2
phosphorus	15	1s22s22p63s23p3
sulfur	16	1s22s22p63s23p4
chlorine	17	1s22s22p63s23p5
argon	18	1s22s22p63s23p6

Order of Filling the Subshells

The following image shows the order for filling the subshells:

     

Determining Electron Configuration

One of the skills you will need to learn to succeed in freshman chemistry is being able to determine the electron configuration of an atom. An electron configuration is basically an account of how many electrons there are, and in what orbitals they reside under "normal" conditions. For example, the element hydrogen (H) has one electron. We know this because its atomic number is one (1), and the atomic number tells you the number of electrons. Where does this electron go? The one electron of hydrogen goes into the lowest energy state it possibly can, which means it will start at "level" one and goes into "s" orbitals first. We say that hydrogen has a "[1s¹]" electron configuration. Looking at the next element on the Periodic Table --helium, or He -- we see it has an atomic number of two, so two electrons. Since " s" orbitals can hold up to two electrons, helium has an electron configuration of "[1s²]".

What about larger atoms? Let's look at carbon, with an atomic number of 6. Where do its 6 electrons go?

• First two: 1s²
• Next two: 2s²
• Last two: 2p²

We can therefore say that carbon has the electron configuration of "[1s²2s²2p²]".

The table below shows the subshells, the number of orbitals, and the maximum number of electrons allowed:

Subshell	Number of Orbitals	Maximum Number of Electrons
s	1	2
p	3	6
d	5	10
f	7	14

The Abridged (shortened) Periodic Table below shows the electron configurations of the elements. Notice for space reasons we sometimes leave off a portion of the electron configuration. For example, look at argon (Ar), element 18. The table below shows its electron configuration as "[3s²3p⁶]" (remembering that "p" orbitals can hold up to six (6) electrons). Its actual electron configuration is:

Ar = [1s²2s²2p⁶3s²3p⁶]

Sometimes you will see the notation: "[Ne]3s²3p⁶", which means to include everything that is in neon (Ne, 10) plus the stuff in the "3"-level orbitals.