In the preceding unit we have learnt that an electron in an atom is characterised by a set of four quantum numbers, and the principal quantum number (n ) defines the main energy level known as shell.
We have also studied about the filling of electrons into different subshells, also referred to as orbitals (s, p, d, f ) in an atom. The distribution of electrons into orbitals of an atom is called its electronic configuration.
An element’s location in the Periodic Table reflects the quantum numbers of the last orbital filled. In this section we will observe a direct connection between the electronic configurations of the elements and the long form of the Periodic Table.
(a) Electronic Configurations in Periods
The period indicates the value of n for the outermost or valence shell. In other words, successive period in the Periodic Table is associated with the filling of the next higher principal energy level (n = 1, n = 2, etc.).
It canbe readily seen that the number of elements in each period is twice the number of atomic orbitals available in the energy level that is being filled.
The first period (n = 1) starts with the filling of the lowest level (1s) and therefore has two elements — hydrogen (ls 1 ) and helium (ls 2 ) when the first shell (K) is completed.
The second period (n = 2) starts with lithium and the third electron enters the 2s orbital. The next element, beryllium has four electrons and has the electronic configuration 1s 2 2s 2 .
Starting from the next element boron, the 2p orbitals are filled with electrons when the L shell is completed at neon (2s 2 2p 6 ). Thus there are 8 elements in the second period.
The third period (n = 3) begins at sodium, and the added electron enters a 3s orbital. Successive filling of 3s and 3p orbitals gives rise to the third period of 8 elements from sodium to argon.
The fourth period (n = 4) starts at potassium, and the added electrons fill up the 4s orbital. Now you may note that before the 4p orbital is filled, filling up of 3d orbitals becomes energetically favourable and we come across the so called 3d transition series of elements.
This starts from scandium (Z = 21) which has the electronic configuration 3d 1 4s 2 . The 3d orbitals are filled at zinc (Z=30) with electronic configuration 3d 104s 2 .
The fourth period ends at krypton with the filling up of the 4p orbitals. Altogether we have 18 elements in this fourth period. The fifth period (n = 5) beginning with rubidium is similar to the fourth period and contains the 4d transition series starting at yttrium (Z = 39).
This period ends at xenon with the filling up of the 5p orbitals. The sixth period (n = 6) contains 32 elements and successive electrons enter 6s, 4f, 5d and 6p orbitals, in the order — filling up of the 4f orbitals begins with cerium (Z = 58) and ends at lutetium (Z = 71) to give the 4f-inner transition series which is called the lanthanoid series.
The seventh period (n = 7) is similar to the sixth period with the successive filling up of the 7s, 5f, 6d and 7p orbitals and includes most of the man-made radioactive elements.
This period will end at the element with atomic number 118 which would belong to the noble gas family. Filling up of the 5f orbitals after actinium (Z = 89) gives the 5f-inner transitionseries known as the actinoid series.
The 4fand 5f-inner transition series of elements are placed separately in the Periodic Table to maintain its structure and to preserve the principle of classification by keeping elements with similar properties in a single column.
(b) Groupwise Electronic Configurations
Elements in the same vertical column or group have similar valence shell electronic configurations, the same number of electrons in the outer orbitals, and similar properties.
For example, the Group 1 elements (alkali metals) all have ns 1 valence shell electronic configuration as shown below.
Thus it can be seen that the properties of an element have periodic dependence upon its atomic number and not on relative atomic mass.