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Ground state electron configurations are the foundation for understanding molecular bonding, properties, and structures. From the electrons in an atom, to the differing orbitals and hybridization, the ground state electron configuration sheds light on many different atomic properties. Fundamentally, understanding electron configuration leads to an understanding of the periodic table.
In 1913, Niels Bohr proposed that electrons could orbit an atom at a certain distance without collapsing into the atom, and that each orbit distance had its own energy level. He proposed that each orbital’s angular momentum, M, was equal to a multiple, n, of Plank’s constant, h, divided by 2π. This gives the equation:
M = nħ where ħ= h/2π and n= 1,2,3,4
This model proposed the Bohr atom, which shows circular orbits surrounding the nucleus.
In addition to having different energy levels, orbitals also have different shapes and orientations, and each can be occupied by two electrons. For each principal quantum number, n, there is one s orbital, three p orbitals, five d orbitals and seven f orbitals. Therefore, an s orbital can hold two electrons, a p orbital can hold six electrons, a d orbital can hold ten electrons, and an f orbital can hold 14 electrons.
There are four quantum numbers n, l, ml, and ms. The principle quantum number n is a positive integer (1,2,3,4) and it represents the energy of the orbital. The angular momentum quantum number l, is from 0 to n – 1. The l values of 0, 1, 2, and 3 correspond to the s, p, d and f orbitals, respectively. The magnetic quantum number ml ranges from –l to +l. This quantum number dictates the orbital orientation, such as px, py, or pz. The quantum spin number ms, is either +1/2 or -1/2 and it dictates the electron spin.
The Aufbau principle states that electrons must fill lowest energy shells first.
Following the model, electrons fill the 1s orbital with two electrons, then the 2s with two electrons, then the 2p with six electrons, then the 3s with two electrons, etc.
There are some exceptions to the Aufbau Principle. This occurs mainly with electrons in the d orbital where extra stability is obtained from a half filled or fully filled d orbital. Therefore, if there are 4 electrons, or 9 electrons in the d orbital, it will move one electron from the s orbital below it to fill the extra space.
|Example 1: Chromium|
|Cr's electron configuration, following the model would be: 1s2 2s2 2p6 3s2 3p6 4s23d4, but instead it is 1s2 2s2 2p6 3s2 3p6 4s13d5, because there is extra stability gained from the half-filled d orbital.|
Hund’s rule states that when filled sub-levels other than s orbital, electrons must not be spin paired in the orbitals until each orbital contains one electron, and no orbital can have two electrons with the same spin (ms).
Pauli Exclusion Principle states that no two electrons can have the same quantum numbers. An orbital can only hold 0, 1, or 2 electrons. They must have opposite spins if there are 2 electrons in the orbittal.
Valence electron shells in the periodic table follow a trend. This can be referred to as the s block, the p block, the d block and the f block (lanthanides and actinides) meaning that, in its ground state, an element in a certain "block" will have its valence electrons in the s, p, d, or f orbitals depending.
Electron configurations are written using the principle quantum number n, followed by the orbital (s, p, d, or f) with the total number of electrons written as a superscript. Example: 1s2 For writing ground state electron configurations, a few main steps should be followed.
Example: Na: 11 e- 1s2 2s2 2p6 3s1 or Na+: 1s2 2s2 2p6
|Example 2: Chromium|
Cr: 1s2 2s2 2p6 3s2 3p64s23d4 half filled orbital, s orbital beneath it
1s2 2s2 2p6 3s2 3p6 4s13d5
Because writing the entire electron configuration can become cumbersome, there is a shorthand option. It is done by using the symbol of the noble gas in the period above the element to represent the electron configuration before it.
Example: Na: [Ne] 3s1
Solution 1. Expanded: 1s2 2s2 2p6 3s2 3p5
Shorthand: [Ne] 3s2 3p5
Solution 2. Expanded: 1s2 2s2 2p6 3s2 3p6 4s13d5
Shorthand: [Ar] 4s13d5
Solution 3. Expanded: 1s2 2s2 2p6 3s2 3p6 4s13d10
Shorthand: [Ar] 4s13d10
Solution 4. Expanded: 1s2 2s2 2p6 3s2 3p6 4s23d5
Shorthand: [Ar} 4s23d5
Solution 5. 1s2 2s2 2p6 3s2 3p6
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