Ground State Electronic Configuration Of Isolated Gaseous Atoms And Ions

You need to be familiar with the structure of an atom. In the structure of an atom electrons are arranged in energy levels or shells, around the nucleus of an atom. As the number of electrons rises, the number of shells also rise around the nucleus. The shell nearest to the nucleus represents the lowest energy level.How many electrons can be filled in each shell? It can be known by the formula 2n². n represents the number of the shell where n is 1,2,3,4 and so on.

Electrons are filled first in lowest energy level and then in higher energy levels, and so on. Each shell further has subshells represented by s, p, d and f. s, p, d and f subshells contain 2, 6, 10 and 14 electrons respectively.

Furthermore, in each subshell, there are orbitals present. A maximum of 2 electrons can be filled in an orbital. Since in each orbital, a maximum,of 2 electrons can be occupied, we can calculate the electron holding capacity of each subshell. We do so by multiplying the number of orbitals in each subshell by 2. In this case, s subshell can have 2 electrons, p subshell can have 6 electrons, d subshell can have 10 electrons and f subshell can have 14 electrons.

Aufbau's principle is based on energy of orbitals or subshells and accordance to the principle that electrons fill lower energy levels first. Filling of electrons happen in increasing order of energy of orbitals. How do we know which orbital has the lowest or highest energy? For this we need to get familiar with the formula (n+ l). Here n is the principle quantum number which represents the number of the shell. l is the orbital quantum number whose value is given by l=n-1. The orbital with least value of n + l will be filled first. We can also predict the increasing order of energy of orbitals, as shown in this illustration. 1s will be filled first due to its lowest energy. Then 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s and so on will be filled according to increasing order of energy of orbitals. Aufbau’s principle does not apply to single electron systems such as the Hydrogen atom, He+ and Li+2 ions.

After understanding the order of energy of orbitals we move towards Hund’s rule. This rule is based on the pairing of electrons. Each orbital can have 2 electrons, as seen in the s orbital. According to this rule, electrons prefer to remain unpaired as much as possible. Electrons in a subshell are paired only when all the orbitals of a subshell are half filled with parallel spin.

Pauli’s exclusion principle tells us about how 2 electrons can be filled correctly in an orbital. According to this principle each orbital can accommodate 2 electrons with opposite spin. Two electrons in an orbital will not have the same set of four quantum numbers.

The spin of electron is represented by the spin quantum number. -1/2 represents spin downwards or counterclockwise spin and +1/2 represents spin upwards or clockwise spin.

Lets take an example of calcium and apply all principles and rules to fill up electrons. The atomic number of Calcium is 20. Therefore, Calcium’s electronic configuration will be 1s²2s²2p⁶3s²3p⁶4s². The first two electrons of calcium will fit in 1s orbital. The next 2 electrons for Calcium fit in the 2s orbital. Then 2p, 3s, 3p and 4s orbitals are filled up applying all principles.

The ascending order of sub energy levels is illustrated as shown.

We got the order of increasing energy level by formula of n+l. Now if we depict this order in ascending order, it can be represented as shown.

Ionization energy is the energy required to remove electron from outermost or valance shell of an atom. The first ionization energy is the energy required to remove the first electron from the valance shell. It is inversely proportional to the radius of an atom. The larger the radius, the smaller will be it’s ionization energy. This is because the valence electrons will be far from the nucleus. Therefore, it will be easy to remove electrons because less energy will be required to do so.

Ionization energy rises along a period of the periodic table. This is because atomic radius lessens along a period of the periodic table as a result of a constant number of shells, and rises in nuclear charge. Helium, Neon and Argon have complete valence shells. Due to this reason, it is difficult to remove electrons from them. In fact, they have the highest ionization energies.

Lithium , Sodium and Potassium have the lowest ionization energies because of the presence of only 1 electron in their outermost shell. Atoms having 1 or 2 electrons in their valance shell are easy to remove. Hence their ionization energy is smaller. All atoms having more electrons in their valance shell need more energy to remove electrons.

Let us discuss the extra stability of d5 and d10 subshells. Either a fully filled or an exactly half filled d subshell is especially stable. To gain this stability, an atom or ion having 1 electron short in d5 or d10 does the following. It shifts an electron from the s subshell of highest energy to the unfilled d subshell. A half filled and fully filled d orbital has the same energy. Therefore these are considered degenerate orbitals.

Due to the same energy of orbitals and symmetrical distribution of electrons the electrons in different orbitals of same subshell exchange their positions. Because of this exchange, some amount of this energy, which we call exchange energy, is released. This makes the atom more stable. So the half filled and fully filled orbitals are more stable.

Their stability is because of exchange energy and symmetry. These orbitals are more symmetrical than any other configuration which leads to greater stability.For example, Chromium is a chemical element with the symbol Cr and atomic number 24. We can write its electronic configuration as 1s²2s²2p⁶3s²3p⁶3d⁴4s². But this electronic configuration is incorrect due to instability of the 3d orbital.

Half-filled and full filled subshells have extra stability. Therefore, one of the 4s² electrons jumps to the 3d⁴ orbital to make it 3d⁵, which is half filled. This gives us the correct configuration as, 1s²2s²2p⁶3s²3p⁶3d⁵4s¹.

Now we shall discuss anomalies of group 2,3,5 and 6. Anomalous behavior is shown by the first member of groups. Every first member differs in properties and often exhibit diagonal relationship with other elements of any other group. In members of group 2. beryllium shows anomalous behavior than other members. It also shows diagonal relationship with aluminum. In members of group 3, boron shows peculiar properties than other elements of the family. It shows diagonal relationship with silicon.

Group 2A of the periodic table are the alkaline earth metals. They are Beryllium, Magnesium, Calcium, Strontium, Barium, and Radium.Group 3A of the periodic table includes the metalloid boron, as well as the metals Aluminum, Gallium, Indium, and Thallium.Group 5A of the periodic table are the pnictogen. They are the nonmetals Nitrogen, and Phosphorus, the metalloids Arsenic and Antimony, and the metal Bismuth.Group 6A of the periodic table are the chalcogens. They are the nonmetals Oxygen, Sulfur, and Selenium, the metalloid Tellurium, and the metal Polonium.

Beryllium shows diagonal relationship with Aluminum as charge to size ratio is the same for both Beryllium and Aluminum. Like Aluminum, Beryllium also does not react with acids due to the presence of oxide layer on metal surface. Like Aluminum, Beryllium oxides can dissolve in excess of alkali as the [Be(OH)4]-2. Both these elements form oxo-anions in strong base. Both of these have bridge bond in their hydrides and chlorides.They have amphoteric oxides. Their oxides are extremely stiff with high melting points. All beryllium compounds and some aluminum compounds have covalent character.

Boron shows diagonal relationship with Silicon. It forms solid acidic oxides like that of silicon. While it is amphoteric in nature, boric acid is a weak acid like that of silicic acid. These have a wide range of polymeric borates and silcates based on shared oxygen atoms between bonding. These both form gaseous oxides. Anomalous behavior of Nitrogen of group 5 shows that nitrogen is gas in nature while others are solids. It has a small size and a high ionization energy which is different in other elements of the group.The non-availability of d-orbital in valance shell and ability to form pi-pi bonds with itself make it anomalous in nature. It is diatomic while others are tetratomic. Due to the absence of a d-orbital, it does not form coordinate bonds.

Now we shall discuss anomalous behavior of oxygen of group 6.Oxygen has a small size than other members of group 6. It has high electronegativity and the absence of a d-orbital in valance shell. It is a non-metal and a gas while others are solid at room temperature. It forms multiple pi-pi bonds with similar size elements. Oxygen is paramagnetic while others are diamagnetic in nature. It forms strong hydrogen bonding in H2O which is not available in H2S.

Anomalous behavior can be due to the following reasons. Elements showing anomalous behavior have a small size as compared to other members of their group. They also have high electronegativity and ionization energy as compared to other members of the group. They have a large charge to radius ratio due to their small size and high ionization energy. They do not have d orbitals in their valence shell. That is why they exhibit anomalous behavior.