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Simple Explanation:
According to the Quantum Mechanical Model of the atom, electrons occupy orbitals, which are regions in space where they are most likely to be found. These orbitals are visualized as electron clouds, with denser areas indicating a higher probability of locating an electron.
As atomic mass increases, atoms contain electrons in higher energy levels. This in turn corresponds to greater distances from the nucleus and higher principal quantum numbers (n).
Electron configuration describes the arrangement of electrons in these orbitals. It also provides understanding of the stability of these electron distributions in atoms and ions.
For a more in-depth explanation of orbitals, click HERE.
Aufbau Principle: Electrons fill the lowest energy orbitals first before moving to higher energy levels.
Notation
Electron configuration is written as a series of numbers and letters:
Numbers represent the principal energy level (n).
Subshells (s, p, d, f) indicate the type of orbital.
Superscripts indicate the number of electrons in that sublevel. Count them by one from left to right.
To write the electron configuration of an atom or ion, start from the 1s orbital, the lowest energy level, and then fill each subshell in order of increasing energy. If a subshell is full, move to the next subshell until all electrons are placed.
How much orbitals can each subshell hold?
S Subshell: 1 orbital (2 electrons)
P Subshell: 3 orbitals (6 electrons)
D Subshell: 5 orbitals (10 electrons)
F Subshell: 7 orbitals (14 electrons)
Sequence of subshells go by their arrangements: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d, 7p.
How many electrons can each energy level hold?
First Energy Level (n=1): 2 electrons
1s
Second Energy Level (n=2): 8 electrons
2s, 2p
Third Energy Level (n=3): 18 electrons
3s, 3p, 3d
Fourth Energy Level (n=4): 32 electrons
4s, 4p, 4d, 4f
Examples of Expanded Electron Configurations
Hydrogen (H): 1s¹
H has 1 electron in 1s.
Magnesium (Mg): 1s² 2s² 2p⁶ 3s²
Mg has 2 electrons in 1s, 2 electrons in 2s, 6 electrons in 2p, and 2 electrons in 3s.
Iron (Fe): 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d⁶
Iodine (I): 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d¹⁰ 4p⁶ 5s² 4d¹⁰ 5p⁵
Tungsten (W): 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d¹⁰ 4p⁶ 5s² 4d¹⁰ 5p⁶ 6s² 4f¹⁴ 5d⁴
UNIQUE: Copper (I) (Cu): 1s², 2s², 2p⁶, 3s², 3p⁶, 4s¹, 3d¹⁰
Due to stability rules, it is more energetically favorable have a full 3d shell and a half 4s shell.
Configuration of Ions:
Sodium Ion (Na⁺): 1s² 2s² 2p⁶
Na⁺ has lost 1 electron, so it has the same configuration as Neon.
Chloride Ion (Cl⁻): 1s² 2s² 2p⁶ 3s² 3p⁶
Cl⁻ has gained 1 electron, filling the 3p orbital.
UNIQUE: Copper (I) Ion (Cu⁺): 1s², 2s², 2p⁶, 3s², 3p⁶, 3d¹⁰
Due to stability rules, it is more energetically favorable to lose an electron in its 4s subshell.
Examples of Noble Gas Configuration Shortcuts
As a shortcut, you could add onto the configurations of Group 18 elements for atoms or ions that surpass those configurations. Basically, you substitute all the configurations that match the previous Group 18 element and move on.
Oxygen (O): [He] 2s² 2p⁴
Phosphorus (P): [Ne] 3s² 3p³
Zinc (Zn): [Ar] 4s² 3d¹⁰
Simple Visualization of Orbital Types
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