Group And Period Wise Electronic Configuration :
The periodic table is a fundamental tool in chemistry, organizing elements based on their atomic number and properties. One key aspect of understanding an element's behaviour is its electronic configuration, which describes the distribution of electrons in its atomic orbitals.
In this article, we delve into the concepts of groupwise and period wise electronic configuration, exploring how they contribute to the periodic trends and properties of elements.
Before delving into groupwise and period wise electronic configuration, it's crucial to grasp the fundamentals of electronic configuration.
The electronic configuration of an atom is represented by a series of numbers and letters, with each element having a unique arrangement of electrons in its orbitals. The three main quantum numbers—n (principal quantum number), l (azimuthal quantum number), and m (magnetic quantum number)—determine the specific arrangement.
Group wise Electronic Configuration :
The periodic table is organized into groups or columns, each containing elements with similar chemical properties. Elements within the same group have identical outer electron configurations, leading to analogous chemical behaviour. For example, elements in Group 1, also known as alkali metals, share the general electronic configuration [noble gas] ns¹.
General outer electronic configuration of elements in groups:
The commonality in group-wise electronic configuration is a consequence of elements within the same group having the same number of valence electrons. Hence, they follow identical patterns.
Period Wise Electronic Configuration :
In contrast to groups, periods in the periodic table are horizontal rows that represent the filling of electron shells. Moving across a period from left to right, the electronic configuration of elements changes incrementally. The period wise electronic configuration highlights the sequential filling of orbitals.
For example, in Period 2, elements exhibit the following electronic configurations: 1s² 2s² for lithium (Li) and 1s² 2s² 2p⁶ for neon (Ne). The shift from one period to the next signifies the occupation of a new principal energy level, leading to variations in atomic size and other properties.
For finding the electronic configuration of the elements in the periods following points are used:
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The period of the element is the value of n, the principal quantum number, for the valence shell.
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The number of electrons that can be accommodated by different energy levels varies.
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The maximum number of electrons that can be accommodated in an energy shell is given by 2n
2
, where n is the energy level. It is the greatest number of electrons that a given energy level can allow. So the first energy level (K shell) can hold up to 2 electrons, the second (L shell) up to 8 electrons, the third (M shell) up to 18 electrons, and so on.
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The second period begins with Lithium and Beryllium, both of which have three and four electrons, respectively, and so the final electrons reach level two.
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The third period begins with Sodium and finishes with Argon, filling the 3s and 3p orbitals in that order. There are eight elements in this period as well.
Periodic Trends and Electronic Configuration :
The arrangement of electrons in the atomic orbitals directly influences periodic trends. For instance, as one moves down a group, the outermost electron occupies a higher energy level, resulting in an increase in atomic size. In contrast, moving across a period, the effective nuclear charge rises, causing a decrease in atomic size.
Similarly, ionization energy and electronegativity are influenced by electronic configuration. Elements with half-filled or completely filled orbitals tend to exhibit higher stability, impacting their ionization energy.
We fill the atomic orbital with the electrons in accordance with these three rules,
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Aufbau Principle
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Pauli Exclusion Principle
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Hund’s Rule
