Electron shells contain sub-shells and orbitals

• In the modern model of the atom, electrons occupy specific energy levels called shells, which surround the nucleus.
• Each shell is assigned an integer number (n = 1, 2, 3...) that indicates its distance from the nucleus, with shells further from the nucleus holding electrons with higher energy.

The maximum number of electrons that can be accommodated in a shell is given by the formula:

2n²

where n is the integer representing the shell.

For example:

• The first shell (n=1) can hold up to 2 x 1² = 2 electrons.
• The second shell (n=2) can hold up to 2 x 2² = 8 electrons.

• Shells are further divided into sub-shells or sublevels, which have distinct energy levels and are labeled as s, p, d, and f.
• Within these sub-shells, electrons are found in orbitals, which are regions of space with a high probability of finding an electron. Each orbital can hold a maximum of two electrons with opposite spins.

The following table summarises the capacity of each sub-shell type to hold electrons:

The distribution of sub-shells across the first four shells is as follows:

Orbitals have characteristic shapes

Orbitals, the regions in which electrons are most likely to be found, have unique shapes:

• Orbitals within the same sub-shell are of equal energy.
• Each orbital can hold two electrons, which must have opposite spins (called spin-pairing).
• s orbitals are spherical in shape
• p orbitals are dumbbell-shaped. The three p orbitals are oriented at right angles to each other.

Representing electron configurations

The arrangement of electrons in an atom is called its electron configuration. This can be depicted in multiple ways:

1. Sub-shell notation

This method uses superscripts to indicate the number of electrons within each sub-shell.

For example, the electron configuration of neon (10 electrons) is 1s^2^ 2s^2^ 2p^6^.

2. Electrons-in-boxes notation

• Orbitals are depicted as boxes, with electrons shown as arrows.
• Oppositely directed arrows represent electrons with opposite spins.
• Electron pairing within orbitals occurs only with opposite spins.

For example, the electron configuration of neon is shown below.

Electron configurations represent the most stable arrangement

Electron configurations are arranged to minimise the overall energy of the atom or ion.

This lowest energy arrangement corresponds to the most stable electronic structure.

To deduce an atom's electron configuration, follow these guidelines:

1. Aufbau principle - Electrons fill the lowest energy orbitals first.

For example, calcium's electron configuration is 1s² 2s² 2p⁶ 3s² 3p⁶ 4s², which can be represented using the electrons-in-boxes notation as follows:

The 4s sub-shell is filled before the 3d sub-shell because the 4s orbital has a lower energy than the 3d orbitals in neutral atoms.

2. Hund's rule - Electrons first occupy orbitals of equal energy singly before pairing up.

3. Pauli exclusion principle - When two electrons occupy the same orbital, they must have opposite spins (up and down) to minimise electron-electron repulsion.

4. For ions in the s and p blocks, electrons are added to or removed from the highest occupied sub-shell, e.g.:

• The electronic configuration of Mg²⁺ is 1s² 2s² 2p⁶_.
• The electronic configuration of Cl⁻ is 1s² 2s² 2p⁶ 3s² 3p⁶.

Note: Noble gas notation can simplify representations, using square brackets for the electron configuration of the preceding noble gas.

For example, calcium's configuration (1s² 2s² 2p⁶ 3s² 3p⁶ 4s²) is abbreviated as [Ar] 4s², where [Ar] represents 1s² 2s² 2p⁶ 3s² 3p⁶.

Transition metals have unusual electron configurations

Chromium (Cr) and copper (Cu) exhibit unusual electron configurations:

• Cr: 1s² 2s² 2p⁶ 3s² 3p⁶ 3d⁵ 4s¹ (instead of 3d⁴ 4s²)
• Cu: 1s² 2s² 2p⁶ 3s² 3p⁶ 3d¹⁰ 4s¹ (rather than 3d⁹ 4s²)

These exceptions occur because configurations with a half-filled (d⁵) or fully filled (d¹⁰) d sub-shell are energetically more favourable.

For transition metals forming ions, the loss of electrons happens from the 4s orbital before the 3d orbital.

For example, consider an iron atom (Fe) with 26 electrons:

• The electron configuration of the iron atom is: 1s² 2s² 2p⁶ 3s² 3p⁶ 3d⁶ 4s²
• When an iron atom loses three electrons to form an Fe³⁺ ion:
• The 4s electrons are lost first.
• Then, one 3d electron is lost.
• The resulting electron configuration of the Fe³⁺ ion with 23 electrons is: 1s² 2s² 2p⁶ 3s² 3p⁶ 3d⁵