Embedded Files
To view the tutorial for writing electron configuration notation, click HERE.
Atomic Orbitals
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 numbers (n).
Introduction to Orbital Diagrams
We could use orbital diagrams as a representation of electron configurations. Here are the basic concepts to it:
Electron Spin: To put it simply, electrons have a property called spin, which create magnetic fields. There are two possible orientations: up-spin (↑) and down-spin (↓).
Pauli Exclusion Principle: In one orbital, there must be two electrons with opposite spins at maximum. No orbital can have two electrons that have the same spin, as that would not be optimal for stability.
Aufbau Principle: Electrons fill the lowest energy orbitals first to maximize attraction to the nucleus.
Hund's Rule: To minimize electron repulsion, electrons must occupy all empty orbitals in a subshell as up-spins before pairing up with down-spins.
FUN FACT: Pairing an up-spin electron with a down-spin electron in an orbital cancels out the magnetism, resulting in diamagnetism. This also means that the more of electrons that possess the same spin in a subshell, the more paramagnetic an atom is.
Up-spins are associated with a north magnetic moment, whereas down-spins are associated with a south magnetic moment.
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
Orbital Diagram Examples
Each box represents an orbital. Boxes packed close together represent a subshell.
General Stability Rules
Full Valence Shell: Atoms or ions are the most stable when the outermost s and p subshells are full, because a full electron configuration is the most energetically favorable.
Examples: Group 18 Elements, Na⁺, Ca²⁺, Br⁻.
Full Subshell: Atoms or ions with full outermost subshells (e.g. ending with 2s², ending with 3p⁶) are more stable than partially-filled subshells, because it maximizes electron pairing while minimizing repulsion.
Examples: Group 18 (full outer p subshell), Group 2 Elements (full outer s subshell).
Half-Filled Subshells: Atoms or ions with half-filled outermost subshells (e.g. ending with 3s¹, ending with 4p³) are more stable than other partially-filled subshells, because all the electrons are symmetrically arranged with the outermost electrons occupying their own orbitals, minimizing repulsion.
Examples: Group 1 (half-filled outer s subshell), Group 15 (half-filled outer p subshell).
Some elements don't follow the standard ordering for electron configuration:
Examples:
Chromium: 1s² 2s² 2p⁶ 3s² 3p⁶ 4s¹ 3d⁵, an electron from 4s goes to 3d to create two half-filled subshells.
Copper: 1s² 2s² 2p⁶ 3s² 3p⁶ 4s¹ 3d¹⁰, an electron from 4s goes to 3d to create a half-filled 4s subshell and a full 3d subshell.
Some elements don't strictly use valence electrons for reactions:
Example:
Iron: It has two valence electrons (config: 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d⁶). However, its most common oxidation state is +3 because it will rather use its two 4s electrons (easiest to lose) and one 3d electron (to create a half 3d subshell).
Titanium: It has two valence electrons (config: 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d²). However, its most common oxidation state is +4 because it will rather use its two 4s electrons (easiest to lose) and two 3d electrons (to empty 3d subshell and expose full 3p subshell).
Page updated
Google Sites
Report abuse