Electronic Configuration of Atoms We know that atomic orbitals are the sub-stationary states or the regions in space where the electrons revolve around the nucleus in an atom. Electronic configuration of atoms is representation of the occupation of electrons in the orbitals. In other words, electronic configuration of atoms specifies the order in which electrons fill up the orbitals in the periodic table. The order in which these electrons are filled into the atomic orbitals are controlled by three crucial guidelines or principles which include: 1. Aufbau or building up principle 2. Hund's rule and 3. Pauli's exclusion principle Let us begin with understanding Aufbau Principle. According to this principle an electron always occupies the lowest energy orbital first before filling the higher level. For example, an electron always occupies 2s, the lower energy orbital, first instead of the higher 3s orbital. The Aufbau or building-up principle can be explained with the example of Hydrogen Atom. Hydrogen has one electron. This electron enters the 1s orbital which has the lowest energy. In other words, building-up principle states that the incoming electrons go to an orbital which has the least (n+l) value. However, the orbital having lower 'n' value will be occupied first, in case any two orbitals have the same (n+l) value. Consider the example of Silicon whose atomic number is 14. Twelve electrons can be accommodated in 1s, 2s, 2p, 3s orbitals. Now, the last two electrons can enter into either 3p or 4s orbital. The (n+ l) values of these orbitals are the same, that is, 3p orbital has a (n+l) value of 3+1=4 and 4s has (n+l) value of 4+0 = 4 This means, both the orbitals have the same (n+l) value. But the 3p orbital has 'n' value, that is 3, which is less than the n value of 4s, which is 4. Therefore thirteenth and fourteenth electrons occupy the 3p orbital first. Thus the electronic configuration of Si is [Ne] 3s 23p2.The superscript represents the number of electrons present in the corresponding orbital. The second important rule to determine the electronic configuration of an atom is the Hund's Rule. It says electron pairing happens only after all the available degenerate orbitals are occupied by one electron each. Let us understand Hund's rule with the help of an example. Consider the element Oxygen with Z=8. I has 8 electrons, the first electron goes into the '1s' orbital of the K-Shell. The second electron will be paired up with the first in the same 1s orbital. Similarly the third and fourth electrons will occupy the 2s orbital of L-Shell. The Fifth electron goes into one of the three 2p orbitals of L-Shell. Let that be 2px. Since the three p-orbitals i.e., 2px,2py and 2pz are degenerate , the sixth electron goes into 2py or 2pz but not 2px. Let us say it goes to 2py. Since 2pz is a degenerate orbital, the seventh electron goes to 2pz instead of pairing up with electron in 2px or 2py. Now, since all the 3 sub-orbitals have one electron each, the eighth electron can pair up with any of the three electrons in 2px, 2py and 2pz orbitals. Thus the electronic configuration of Oxygen can be written as 1s1.2s2. 2px2 .2py1. 2pz1. The arrows indicate electrons with spin +1/2 and -1/2. Let us consider the nitrogen atom. it has 7 electrons. The first six electrons have the same arrangement as that of carbon atom 1s1.2s2. 2px1 .2py1. The seventh electron will enter only in 2pz but can not enter into 2px or 2py orbital. Thus the configuration is 1s1.2s2. 2px1 .2py12pz1. The third important rule for electronic configuration, that is, the Pauli's Exclusion Principle states that no two electrons will have the four quantum numbers same. This means that two electrons can ever have any identical values of n, l, m and s values. Because of this rule a single orbital can have only 2 electrons.