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1.2 Electronic Structure and Periodic Properties of Elements

Learning Objectives

By the end of this section, you will be able to:

  • Derive the predicted ground-state electron configurations of atoms
  • Identify and explain exceptions to predicted electron configurations for atoms and ions
  • Relate electron configurations to element classifications in the periodic table

Electronic configuration is a fundamental concept in the field of chemistry and plays a pivotal role in understanding the behaviour and properties of atoms and molecules. It provides a systematic way of representing the arrangement of electrons within an atom, describing how they are distributed in various energy levels and orbitals. This arrangement is critical because it directly influences an element’s chemical reactivity, bonding capabilities, and its position on the periodic table.

Atoms are composed of three fundamental subatomic particles: protons, neutrons, and electrons. The electronic configuration is a precise and organised representation of the distribution of these electrons within an atom’s various orbitals. An electron’s “address” within an atom is composed of successively narrower categories – shells, subshell, and orbital, which are determined by quantum mechanics.

Electron shells are the primary energy levels where electrons are found. The shell is designated using the letter n (n = 1, 2, 3, …) and is organised by increasing energy. The first shell (n = 1) is the closest to the nucleus and can hold a maximum of 2 electrons, while the second shell (n = 2) can hold up to 8 electrons. The third shell (n = 3) can accommodate a maximum of 18 electrons, and this pattern continues as one moves down the periodic table. The further the shell from the nucleus, the higher the energy of its electrons.

Within each energy level, there are different types of subshells, which are labelled using letters (s, p, d, f). The first shell has only one subshell, s.  The second shell has two subshells, s and p. The third shell has three subshells, s, p, and d. The fourth shell has four subshells, s, p, d, and f.

Finally, within each subshell, electrons are grouped into orbitals, which are the regions within subshells where electrons are most likely to be found. An s subshell has only one orbital, while p, d, and f subshells have three, five, and seven orbitals, respectively. Each orbital can hold maximum of two electrons with opposite spin, a property governed by quantum mechanics. This means the s subshell can hold a maximum of 2 electrons, the p subshell can hold up to 6 electrons, the d subshell can hold 10 electrons, and the f subshell can hold 14 electrons. Table 1.2.1 summarises the distribution of electrons in atoms.

Table 1.2.1 Electron Distribution in Atoms

Shell number 1 2 3 4
Subshell designation s s, p s, p, d s, p, d, f
Number of orbitals 1 1, 3 1, 3, 5 1, 3, 5, 7
Maximum number of electrons 2 2, 6 2, 6, 10 2, 6, 10, 14
Total electron in shell 2 8 18 32

We describe an electron configuration with a symbol that contains three pieces of information (Figure 1.2.1):

  1. The number of  shell, n,
  2. The letter that designates the subshell (s, p, d, f), and
  3. A superscript number that designates the number of electrons in that particular subshell.

The Electron Configurations

 

The arrangement of electrons in an atom can be predicted by applying three rules:

Rule 1: Electrons occupy the lowest-energy orbitals available (Aufbau Principle)

The order of orbital energy levels is shown in Figure 1.2.2

Rule 2: Each orbital can hold up to two electrons (Pauli Exclusion Principle)

Rule 3: Electrons will occupy degenerate orbitals (orbitals of equal energy) singly before pairing up, maximising the number of unpaired electrons (Hund’s Rule)

This figure includes a chart used to order the filling of electrons into atoms. At the top is a blue circle labeled “1 s.” In a row beneath this circle are 6 additional blue circles labeled “2 s” through “7 s.” A column to the right begins just right of 2 s and contains pink circles labeled 2 p through 7 p. A column to the right begins just right of 3 p and contains yellow circles labeled 3 d through 6 d. No circles are placed to the right of the 7 s and 7 p circles. A final column on the right begins right of 4 d. It includes grey circles labeled, “4 f” and, “5 f.” No circles are placed right of 6 d. Through these circles, arrows are included in the figure pointing down and to the left. The first arrow begins in the upper right and passes through 1 s. The second arrow begins just below and passes through 2 s. The third arrow passes through 2 p and 3 s. The fourth arrow passes through 3 p and 4 s. This pattern of parallel arrows pointing downward to the left continues through all circles completing the pattern 1 s 2 s 2 p 3 s 3 p 4 s 3 d 4 p 5 s 4 d 5 p 6 s 4 f 5 d 6 p 7 s 5 f 6 d 7 p.
Figure 1.2.2 This diagram depicts the energy order for atomic orbitals and is useful for deriving ground-state electron configurations.
In this figure, a periodic table is shown that is entitled, “Electron Configuration Table.” Beneath the table, a square for the element hydrogen is shown enlarged to provide detail. The element symbol, H, is placed in the upper left corner. In the upper right is the number of electrons, 1. The lower central portion of the element square contains the subshell, 1s. Helium and elements in groups 1 and 2 are shaded blue. In this region, the rows are labeled 1s through 7s moving down the table. Groups 3 through 12 are shaded orange, and the rows are labeled 3d through 6d moving down the table. Groups 13 through 18, except helium, are shaded pink and are labeled 2p through 6p moving down the table.
Figure 1.2.3 This partial periodic table shows electron configurations for the valence subshells of atoms. By “building up” from hydrogen, this table can be used to determine the electron configuration for atoms of most elements in the periodic table.

We will now construct the electron configuration and orbital diagram for a selection of atoms in the first and second periods of the periodic table. Orbital diagrams are pictorial representations of the electron configuration, showing the individual orbitals and the pairing arrangement of electrons. We start with a single hydrogen atom (atomic number 1), which consists of one proton and one electron. The electron configuration and the orbital diagram are:

In this figure, the element symbol H is followed by the electron configuration is 1 s superscript 1. An orbital diagram is provided that consists of a single square. The square is labeled below as, “1 s.” It contains a single upward pointing half arrow.

Following hydrogen is the noble gas helium, which has an atomic number of 2. The helium atom contains two protons and two electrons. For orbital diagrams, this means two arrows go in each box (representing two electrons in each orbital) and the arrows must point in opposite directions (representing paired spins). The electron configuration and orbital diagram of helium are:

In this figure, the element symbol H e is followed by the electron configuration, “1 s superscript 2.” An orbital diagram is provided that consists of a single square. The square is labeled below as “1 s.” It contains a pair of half arrows: one pointing up and the other down.

The n = 1 shell is completely filled in a helium atom.

The next atom is the alkali metal lithium with an atomic number of 3. The first two electrons in lithium fill the 1s orbital. The remaining electron must occupy the orbital of next lowest energy, the 2s orbital. Thus, the electron configuration and orbital diagram of lithium are:

In this figure, the element symbol L i is followed by the electron configuration, “1 s superscript 2 2 s superscript 1.” An orbital diagram is provided that consists of two individual squares. The first square is labeled below as, “1 s.” The second square is similarly labeled, “2 s.” The first square contains a pair of half arrows: one pointing up and the other down. The second square contains a single upward pointing arrow.

Carbon (atomic number 6) has six electrons. Four of them fill the 1s and 2s orbitals. The remaining two electrons occupy the 2p subshell. We now have a choice of filling one of the 2p orbitals and pairing the electrons or of leaving the electrons unpaired in two different, but degenerate, p orbitals. The orbitals are filled as described by Hund’s rule: the lowest-energy configuration for an atom with electrons within a set of degenerate orbitals is that having the maximum number of unpaired electrons. The electron configuration and orbital diagram for carbon are:

In this figure, the element symbol C is followed by the electron configuration, “1 s superscript 2 2 s superscript 2 2 p superscript 2.” The orbital diagram consists of two individual squares followed by 3 connected squares in a single row. The first blue square is labeled below as, “1 s.” The second is similarly labeled, “2 s.” The connected squares are labeled below as, “2 p.” All squares not connected to each other contain a pair of half arrows: one pointing up and the other down. The first two squares in the group of 3 each contain a single upward pointing arrow.

Following the same rules, the electron configurations and orbital diagrams of N, O, F, and Ne are:

This figure includes electron configurations and orbital diagrams for four elements, N, O, F, and N e. Each diagram consists of two individual squares followed by 3 connected squares in a single row. The first square is labeled below as, “1 s.” The second is similarly labeled, “2 s.” The connected squares are labeled below as, “2 p.” All squares not connected to each other contain a pair of half arrows: one pointing up and the other down. For the element N, the electron configuration is 1 s superscript 2 2 s superscript 2 2 p superscript 3. Each of the squares in the group of 3 contains a single upward pointing arrow for this element. For the element O, the electron configuration is 1 s superscript 2 2 s superscript 2 2 p superscript 4. The first square in the group of 3 contains a pair of arrows and the last two squares contain single upward pointing arrows. For the element F, the electron configuration is 1 s superscript 2 2 s superscript 2 2 p superscript 5. The first two squares in the group of 3 each contain a pair of arrows and the last square contains a single upward pointing arrow. For the element N e, the electron configuration is 1 s superscript 2 2 s superscript 2 2 p superscript 6. The squares in the group of 3 each contains a pair of arrows.

The alkali metal sodium (atomic number 11) has one more electron than the neon atom. This electron must go into the lowest-energy subshell available, the 3s orbital, giving a 1s22s22p63s1 configuration. The electrons occupying the outermost shell orbital(s) (highest value of n) are called valence electrons, and those occupying the inner shell orbitals are called core electrons (Figure 1.2.4). Since the core electron shells correspond to noble gas electron configurations, we can abbreviate electron configurations by writing the noble gas that matches the core electron configuration, along with the valence electrons in a condensed format. For our sodium example, the symbol [Ne] represents core electrons, (1s22s22p6) and our abbreviated or condensed configuration is [Ne]3s1.

This figure includes the element symbol N a, followed by the electron configuration for the element. The first part of the electron configuration, 1 s superscript 2 2 s superscript 2 2 p superscript 6, is shaded in purple and is labeled, “core electrons.” The last portion, 3 s superscript 1, is shaded orange and is labeled, “valence electron.” To the right of this configuration is the word “Abbreviation” followed by [ N e ] 3 s superscript 1.
Figure 1.2.4 A core-abbreviated electron configuration (right) replaces the core electrons with the noble gas symbol whose configuration matches the core electron configuration of the other element.
The electron configurations of silicon (14 electrons), phosphorus (15 electrons), sulfur (16 electrons), chlorine (17 electrons), and argon (18 electrons) are analogous in the electron configurations of their outer shells to their corresponding family members carbon, nitrogen, oxygen, fluorine, and neon, respectively, except that the principal quantum number of the outer shell of the heavier elements has increased by one to n = 3. Figure 1.2.5 shows the lowest energy, or ground-state, electron configuration for these elements as well as that for atoms of each of the known elements.
A periodic table, entitled, “Electron Configuration Table” is shown. The table includes the outer electron configuration information, atomic numbers, and element symbols for all elements. A square for the element hydrogen is pulled out beneath the table to provide detail. The blue shaded square includes the atomic number in the upper left corner, which is 1; the element symbol, H, in the upper right corner; and the outer electron configuration in the lower, central portion of the square. For H, this is 1s superscript 1.
Figure 1.2.5 This version of the periodic table shows the outer-shell electron configuration of each element. Note that down each group, the configuration is often similar.

Check Your Learning

Identify the atoms from the electron configurations given:

(a) [Ar]4s23d5

(b) [Kr]5s24d105p6

ANSWER:

(a) Mn (b) Xe

Electron Configurations and the Periodic Table

 

As described earlier, the periodic table arranges atoms based on increasing atomic number so that elements with the same chemical properties recur periodically. When their electron configurations are added to the table (Figure 1.2.5), we also see a periodic recurrence of similar electron configurations in the outer shells of these elements. Because they are in the outer shells of an atom, valence electrons play the most important role in chemical reactions. The outer electrons have the highest energy of the electrons in an atom and are more easily lost or shared than the core electrons. Valence electrons are also the determining factor in some physical properties of the elements.

Elements in any one group (or column) have the same number of valence electrons; the alkali metals lithium and sodium each have only one valence electron, the alkaline earth metals beryllium and magnesium each have two, and the halogens fluorine and chlorine each have seven valence electrons. The similarity in chemical properties among elements of the same group occurs because they have the same number of valence electrons. It is the loss, gain, or sharing of valence electrons that defines how elements react.

It is important to remember that the periodic table was developed on the basis of the chemical behaviour of the elements, well before any idea of their atomic structure was available. Now we can understand why the periodic table has the arrangement it has—the arrangement puts elements whose atoms have the same number of valence electrons in the same group. This arrangement is emphasized in Figure 1.2.5, which shows in periodic-table form the electron configuration of the last subshell to be filled by the Aufbau principle. The colored sections of Figure 1.2.5 show the three categories of elements classified by the orbitals being filled: main group, transition, and inner transition elements. These classifications determine which orbitals are counted in the valence shell, or highest energy level orbitals of an atom.

  1. Main group elements (sometimes called representative elements) are those in which the last electron added enters an s or a p orbital in the outermost shell, shown in blue and red in Figure 1.2.5. This category includes all the nonmetallic elements, as well as many metals and the metalloids. The valence electrons for main group elements are those with the highest n level. For example, gallium (Ga, atomic number 31) has the electron configuration [Ar]4s23d104p1, which contains three valence electrons (underlined). The completely filled d orbitals count as core, not valence, electrons.
  2. Transition elements or transition metals. These are metallic elements in which the last electron added enters a d orbital.
  3. Inner transition elements are metallic elements in which the last electron added occupies an f orbital. They are shown in green in Figure 1.2.5.

Section Summary

  • In an atom, electrons are distributed in narrower categories – shells, subshell, and orbital.
  • The way electrons is arranged directly influences an element’s chemical reactivity and bonding capabilities.
  • The electronic configuration can be predicted using three rules: the Aufbau Principle, the Pauli Exclusion Principle, and Hund’s Rule.

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