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Unit 4: Organizing Atoms and Electrons—The Periodic Table

Section 6: Electron Configurations

Keeping track of all the electrons in an atom can be an intimidating task. In order to organize and inventory where all the electrons in an atom are, we use something called the "electron configuration." The placement of electrons in an atom dictates how each atom behaves, what compounds it will form, and how reactive it is. For example, full sets of orbitals confer stability, and unpaired electrons are highly reactive. In order to get to the point where we can understand and use electron configurations, first we need to understand some of the basic rules that electrons follow in orbitals.

Orbital Blocks on the Periodic Table

Figure 4-8. Orbital Blocks on the Periodic Table

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Orbital Blocks on the Periodic Table

Figure 4-8. Orbital Blocks on the Periodic Table

The periodic table can be divided into blocks based on the location of the outermost electrons in the atom—called "valence electrons." The first two columns are the s block. Columns 11–18 are the p block, 3–10 are the d block, and the bottom two rows are the f block. The blocks, in addition to the period or row numbers, can be used to determine the electron configuration of any element.

Aufbau Principle

The Aufbau Principle states that electrons will fill the lowest energy orbital first. The word Aufbau is the German word meaning "filling." Roughly, the orbitals that are closest in proximity to the nucleus are the lowest in energy. Every atom starts by filling the 1s orbital first. However, the actual order of the orbitals is empirical, which means it must be experimentally determined. So how do we know the order of all these orbitals? Conveniently, the periodic table is laid out in this order from left to right, as can be seen in Figure 4-8. The order of the subshell energies can be transcribed from the periodic table and is represented as a list here:

1s 2s 2p 3s 3p 4s 3d 4p 5s 4d 5p 6s 4f 5d 6p 7s

As we can see, the layout and design of the periodic table provide a shortcut for filling orbitals based on the rows and orbital blocks. The first two columns of the periodic table are known as the s block; the middle columns are the d block; the bottom rows are the f block; and the far right columns are the p block. These four letters represent the outermost subshells of each element. For the main group elements, each row is a new shell. For example, the third row is the beginning of the third shell, which contains a 3s and a set of 3p orbitals. Or if we look in the sixth row, we'll see these elements are filling the 6s and 6p. The d and f blocks do not quite follow the shell pattern. This is just part of the quirk of nature that all of the subshells in a shell are not at the exact same energy, and sometimes orbitals from a different shell can be lower in energy. For example, 3d is higher in energy than 4s.

Hund's Rule
Hund's Rule

Figure 4-9. Hund's Rule

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Hund's Rule

Figure 4-9. Hund's Rule

In this example for nitrogen, the electrons in the 2p subshell should be placed in each of the three available orbitals before doubling up. Electrons repel each other, and this method accounts for electrons naturally maximizing the space between themselves.

A German physicist, Friedrich Hund (1896–1997), developed a rule for filling a set of orbitals that all have equal energies. The p, d, and f subshells have multiple orbitals at the same energy. Every time there is a p subshell, it comes as a set of three orbitals, each of which can hold a maximum of two electrons. In Figure 4-6 of Section 5 of this Unit, lines represent orbitals. Note that 2p, 3p, and 4p subshells are drawn as three orbitals at the same energy level. Any d or f orbital will be drawn with 5 or 7 orbitals, respectively.

Hund's Rule can also be thought of as the Bus Rule: Imagine getting onto a city bus. There are many empty double seats. They slowly fill up with one person in each. Riders generally choose empty double seats before sitting down next to a stranger. Just like an electron, wouldn't we sit in the empty seat?

Pauli Exclusion Principle

In orbital notation, arrows pointing up or down represent electrons. An arrow that points up is representing a "spin up" electron, and an arrow that points down represents a "spin down" electron. There is another fundamental property of subatomic particles that is called "spin." However, nothing is actually spinning on the electron; that's just the name of this property. An electron can have only two possible spin values, up or down.

Electrons in an atom must follow the Pauli Exclusion Principle, which states that multiple electrons in the same orbital cannot have the same spin. Since there are only two types of spin, each orbital can only hold, at most, two electrons (one that is spin up and one that is spin down). In Figure 4-9, the three electrons in the 2p orbital are all of the same type of spin, one in each orbital. The next electron would be pointed down to represent the other type of spin, and placed in any one of the 2p orbitals to avoid violating the Pauli Exclusion Principle.

In summary, there are three basic principles for arranging electrons in an atom. First, fill orbitals from lowest energy to highest; second, put one electron into each orbital of a subshell before pairing up electrons in the same orbital; and third, place, at most, two electrons, with opposite spins in any given orbital. The electron configurations of the first 12 elements are shown in Table 4-3.

Table 4-3. Electron Configurations for the First 12 Elements
ElementElectron ConfigurationElementElectron Configuration
Noble Gas Configuration

As a shortcut, we can use noble gas configuration. The noble gases represent a full shell of electrons, which is why they do not react readily with other elements. Instead of writing out 1s22s22p63s2 for magnesium, we can replace 1s22s22p6 with [Ne]. The noble gas configuration for magnesium would be [Ne]3s2. The inner electrons represented by [Ne] are lower in energy and more stable than the two electrons in the 3s2 orbitals. Those electrons would be the ones involved in reactions.

Valence Electrons and Metals

In Section 5, we learned that valence electrons are the electrons that dictate atomic behavior. For the main group and non-metallic elements, it is easy to tell how many valence electrons an element has just by looking at the periodic table. As we move from left to right, each column has one more valence electron. By the time we reach the final column of noble gases, all the orbitals are full.

However, there are a lot of metallic elements on the table, and their valence electrons are a bit more complicated. Because they are what we call f-block or d-block elements, their valence electrons are actually a combination of their highest-level s orbital as well as some other electrons. The d-block elements, which we call the "transition metals," have their valence electrons in both the s orbital and the d orbitals. The f-block elements, the lanthanides and actinides, have their valence electrons in their outermost s orbital and in their f orbitals as well. The chemistry of metals is very complicated and will be revisited in Units 11 and 13.



A quantum mechanical property associated with an electron that can be plus or minus one-half.


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