Lecture notes
Electronic configuration
- Module
- 31627H
- Institution
- PEARSON (PEARSON)
This document is an indepth explination of Electronic configuration. It goes over everything you would need to know in order to pass
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Electronic structure of Atoms for BTEC studentsElectronic configuration
Electronic configuration refers to the arrangement of electrons in an atom, ion, or molecule.
It provides a way to describe the distribution of electrons across different energy levels,
orbitals, and subshells within an atom or molecule.
Atoms, ions or molecules have a nucleus, which contains protons and neutrons, and
electrons that orbit the nucleus in specific energy levels called shells. Each shell can hold a
certain number of electrons. The first shell can hold up to 2 electrons maximum. The second
shell can hold up to 8 electrons maximum. The third shell can hold up to 18 electrons
maximum, and so on.
Within each shell, there are different subshells or orbitals that have different shapes and
energy levels.
The first shell only has one subshell, called the s orbital, which can hold a maximum of 2
electrons.
The second shell has two subshells: the s orbital and the p orbital. The s orbital can hold 2
electrons, while the p orbital can hold 6 electrons.
The third shell has three subshells: the s, p, and d orbitals. The s orbital can still hold 2
electrons, the p orbital can hold 6 electrons, and the d orbital can hold 10 electrons.
The fourth shell has four subshells: the s, p, d and f orbitals. The s orbital can still hold 2
electrons, the p orbital can hold 6 electrons, the d orbital can hold 10 electrons, and the f
orbital can hold 14 electrons
The order in which the different subshells are filled follows a specific pattern called the
Aufbau principle. According to this principle, lower energy orbitals fill up before higher energy
orbitals. The order of filling is 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f,
6d, 7p, 8s and so on.
To represent the electronic configuration of an atom, the subshells and the number of
electrons in each subshell are written using a standardised notation.
For example, the electronic configuration of carbon is 1s2, 2s2, 2p,2. This indicates that
carbon has 2 electrons in the 1s subshell, 2 electrons in the 2s subshell, and 2 electrons in
the 2p subshell.
There are a few exceptions to the normal electronic configuration patterns in certain
elements due to the nature of electron filling in atomic orbitals. These exceptions occur to
achieve more stable half-filled or fully-filled sublevels.
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