Understanding electron configurations of s- and p-block ions is essential for predicting chemical behaviour, charge formation, periodic trends, and the stability of the resulting electronic structures.

Electron Configurations of s- and p-Block Ions

Electron configurations describe how electrons are distributed across atomic orbitals, and this arrangement governs the formation of ions in the s- and p-block. Ions form when atoms lose or gain electrons to achieve more stable electronic structures, often resembling noble gas configurations. For OCR A-Level Chemistry, students must be able to deduce electron configurations for ions with atomic numbers up to 36, using sub-shell notation and understanding the effect of ionic charge on electron number.

The Basis of Ion Formation in the s- and p-Block

s- and p-block elements form ions through predictable patterns related to their group number.

• Group 1 and 2 metals (s-block): lose electrons to form cations.

• Groups 15–17 non-metals (p-block): gain electrons to form anions.

• Group 18 elements: virtually never form ions due to full outer shells.

Electrons are always removed or added from/to the highest-energy occupied sub-shell, which corresponds to the outermost shell for s- and p-block elements.

Understanding Electron Loss in s-Block Cations

s-block metals form ions by losing electrons from their outermost s orbital. This process creates cations with reduced electron counts but unchanged proton numbers.

When the term cation is first introduced, it requires a definition.

After an electron configuration is determined, the cation’s configuration must reflect the removal of electrons from the highest occupied shell, not the order of sub-shell filling.

Normal electrons in the s-block are removed as follows:

• Loss of one electron → Group 1 elements form 1+ ions.

• Loss of two electrons → Group 2 elements form 2+ ions.

Key points for s-block ions:

• Electrons are removed from the ns sub-shell first.

• No d or f orbitals are involved for elements up to Z = 36.

Electron Gain in p-Block Anions

Non-metals in the p-block typically form anions by gaining electrons to fill their p sub-shells. These elements move towards the stable configuration of the nearest noble gas.

When the term anion is first introduced, it requires a definition.

Students must be able to assign electrons to the 2p, 3p, and 4p sub-shells appropriately, depending on atomic number and the ionic charge.

General patterns include:

• Group 17 elements gain one electron → 1− ions.

• Group 16 elements gain two electrons → 2− ions.

• Group 15 elements gain three electrons → 3− ions (less common but examinable).

Using Sub-Shell Notation for Ion Configurations

Electron configurations for ions must always be written in sub-shell notation, e.g. 2s² 2p⁶, not in box notation or dot notation.

• Identify the atomic number to determine the number of protons.

• Calculate the electron number in the ion:

  • For cations → subtract the ionic charge.

  • For anions → add the ionic charge.

• Write the neutral atom configuration first using the standard filling order (1s → 2s → 2p → 3s → 3p → 4s → 3d → 4p).

Caption: This diagram shows the order in which electrons fill atomic orbitals, illustrating energy progression and the placement of s, p, d, and f subshells. It helps visualise why electrons occupy orbitals in this sequence when determining electron configurations. Some subshells shown extend beyond the s- and p-block requirements for this topic. Source

• Adjust the configuration by removing or adding electrons strictly from/to the highest principal quantum number (n).

This adjustment is critical; although 4s fills before 3d, electrons are removed from 4s before 3d only in transition metals — but this does not apply here because the sub-topic is restricted to s- and p-block ions.

Important Principles for Determining Ion Configurations

s- and p-block ions exhibit predictable behaviour. Students should internalise the following facts:

• The outermost shell (highest n value) is always where electron addition or removal occurs.

• No d electrons are involved for any element ≤ Z = 36 when forming simple s- or p-block ions.

• A stable configuration often matches the nearest noble gas (He, Ne, Ar, Kr).

• The effective number of electrons in the highest-energy p sub-shell governs the reactivity and ion formation tendencies.

Recognising Stability in Ion Configurations

The concept of stability is linked to achieving full s and p sub-shells. When an ion reaches configurations such as:

• ns² np⁶ (full outer shell), or

• 1s² for hydrogen or lithium ions,
  it attains a structure that reflects noble gas stability.

When the term noble gas configuration appears for the first time, a definition is necessary.

Visualising Electron Redistribution

Although diagrams are not required in these notes, students should mentally picture how electrons shift between shells and sub-shells. Clear visualisation reinforces understanding of:

• The sequential structure of shells: n = 1 to n = 4.

• The boundaries of s and p sub-shell capacities:

  • s sub-shell → 1 orbital → 2 electrons.

  • p sub-shell → 3 orbitals → 6 electrons.

• The directional flow of electrons (loss or gain) during ion formation.

This diagram highlights the s- and p-block regions of the periodic table, supporting identification of which elements form the ions discussed in this topic. The d- and f-blocks appear but are not required for the OCR specification on s- and p-block ion configurations. Source

These patterns enable quick deduction of correct ion configurations when ionic charges and atomic numbers are provided.

This diagram illustrates electron transfer between sodium and chlorine to form Na⁺ and Cl⁻, showing how electron gain and loss create stable noble-gas-like configurations. Although the figure also introduces ionic bonding, only the electron-transfer aspect is relevant to this subtopic. Source