1.3 Electrons, energy levels and atomic orbitals

Use this electron-structure sequence for ground-state atoms and ions within the assessed range hydrogen to krypton. The scope tells you which cases to handle; it is not a new chemical rule.

Within that range, describe electrons using principal shells, s/p/d/f sub-shells and orbitals, and use full or noble-gas shorthand configurations as required. Fill lower-energy sub-shells first, while keeping the stated 4s/3d ordering and any supported configuration exception in view.

For an ion, adjust the electron count for the charge and then write the configuration. For transition-metal ions, remove 4s electrons before 3d electrons; do not apply the neutral-atom filling order mechanically when forming the ion.

Do not import excited-state arrangements or examples beyond krypton into a scope-limited answer. Conversely, an out-of-scope example is not evidence that a broader chemical statement is false; it only lies outside the cases this card assesses. Ionisation-energy trends belong to the next Level-3 topic.

Use the hierarchy in order: a principal shell is an energy level labelled by n; each shell contains sub-shells labelled s, p, d and, where supported, f; each sub-shell contains atomic orbitals. An orbital is an allowed region, not a fixed circular path.

Each orbital holds at most two electrons. Therefore s has 1 orbital and holds 2 electrons, p has 3 orbitals and holds 6, d has 5 and holds 10, and f has 7 and holds 14. For example, 2p means the p sub-shell in shell n = 2.

Higher principal shells are generally higher in energy, while sub-shell energy follows the supported order s < p < d < f with overlap at higher shells, including 4s/3d. s orbitals are spherical; p sub-shells contain three perpendicular dumbbell-shaped orbitals, pₓ, pᵧ and p_z.

Do not confuse shell, sub-shell and orbital, or treat an orbital as a track around the nucleus. The detailed filling and electron-count method belongs to the neighbouring electron-configuration objective; this card supplies the location, capacity and model boundaries.

A sub-shell is made of atomic orbitals: an s sub-shell contains 1 orbital, a p sub-shell contains 3 orbitals, and a d sub-shell contains 5 orbitals. The three p orbitals are conventionally labelled pₓ, pᵧ and p_z.

Each orbital can hold a maximum of 2 electrons. Therefore the corresponding sub-shell capacities are s = 1 × 2 = 2 electrons, p = 3 × 2 = 6 electrons and d = 5 × 2 = 10 electrons. Where included, f has 7 orbitals and a capacity of 14 electrons.

To determine a sub-shell capacity, identify its orbital count and multiply by two. In a ground-state p sub-shell, orbitals of the same sub-shell have equal energy; electrons occupy separate orbitals before pairing, subject to the electron-configuration rules.

Do not confuse the number of orbitals with the number of electrons, and do not infer a new shell from the s/p/d label alone. Detailed filling sequences, ions and exceptions belong to the neighbouring electron-configuration objective; detailed d-orbital shape is not required at AS Level.

Build a ground-state configuration by filling available sub-shells from lower to higher energy. Shell number alone does not give the complete order because sub-shell energies overlap at higher levels; in particular, 4s is filled before 3d for a neutral atom.

Use a fixed method: count the electrons, follow the supported increasing-energy sequence, place no more than the allowed number in each sub-shell, and check that the exponents total the required electron count. A shorthand configuration uses a preceding noble-gas core where appropriate.

For potassium, the argon core is followed by 4s¹, so its shorthand configuration is [Ar] 4s¹ rather than [Ar] 3d¹. The 4s/3d overlap is the reason the simple numerical shell order is not a safe filling rule.

Do not transfer neutral-atom filling order unchanged to a transition-metal ion: when such ions form, electrons are removed from 4s before 3d. Keep filling order, ion formation and special exceptions distinct; orbital capacities are covered by the neighbouring sub-shell objective.

Read an electron configuration from left to right: the leading number identifies the principal shell, the letter identifies the sub-shell, and the superscript gives the number of electrons in that sub-shell. For example, 2p⁴ means four electrons in the p sub-shell of shell n = 2.

Determine the species’ electron count before checking the notation: a neutral atom has electrons equal to its atomic number, a positive ion has lost electrons, and a negative ion has gained electrons. The superscripts in a complete configuration must add to that count.

A full configuration lists occupied sub-shells from the beginning; shorthand replaces the inner-electron part with the nearest noble-gas symbol in brackets. The outer occupied sub-shells identify the valence-electron region relevant to the species’ outer electronic structure.

Use the notation to report electron arrangement, not as a substitute for the separate filling-order rules. Check charge and total superscripts before accepting an ion configuration, and do not confuse a sub-shell superscript with an orbital count or with the atomic number.

In a ground-state sub-shell, orbitals of the same sub-shell have equal energy. Electrons occupy these orbitals singly before any pairing, and the single electrons have parallel spins. This is Hund’s rule.

Once every equivalent orbital contains one electron, additional electrons pair in the orbitals. The two electrons in a paired orbital must have opposite spins; this arrangement limits the effect of electron–electron repulsion within the sub-shell.

Apply the rule by drawing one box per orbital: p³ is ↑, ↑, ↑, whereas p⁴ is ↑↓, ↑, ↑. The arrows show spin and the number of unpaired electrons, not merely the total electron count.

Do not pair electrons in the first orbital while equivalent orbitals remain empty, and do not give paired electrons the same spin. Keep Hund’s rule for equal-energy orbitals distinct from the separate order in which sub-shells are filled.

Use this sequence for a ground-state configuration: (1) determine the species’ electron count, (2) fill sub-shells in increasing energy, (3) respect each sub-shell’s orbital capacity and Hund’s rule, and (4) add the superscripts to check the total. Use full or noble-gas shorthand notation as requested.

For an ion, start from the neutral atom’s proton number and adjust only the electron count for the charge. A positive ion has fewer electrons and a negative ion has more; the proton number and element identity do not change.

When forming a transition-metal ion, remove electrons from 4s before 3d, even though 4s was filled before 3d in the neutral atom. For example, Fe is [Ar] 4s² 3d⁶, whereas Fe²⁺ is [Ar] 3d⁶.

Reject a configuration if its superscripts do not match the species’ electron count or if it violates capacity, filling order or spin occupancy. Do not remove 3d before 4s or treat ion formation as a proton-number change; special Cr/Cu arrangements remain separate supported exceptions.

In electron-box notation, each box represents one atomic orbital and each arrow represents one electron. Opposite arrow directions in one box show a pair with opposite spins.

Draw the required boxes for the sub-shell, place electrons in the supported energy order, and use one arrow in each equivalent orbital before pairing. The diagram therefore displays both the number of electrons and their orbital occupancy.

For a 2p⁴ sub-shell, draw three boxes as ↑↓, ↑, ↑. There are four arrows in total, one paired box and two unpaired electrons; the arrow directions make the spin information visible.

A box diagram and superscript notation represent the same configuration in different formats. Do not pair electrons while an equivalent orbital is empty, and do not interpret a box as a shell or a fixed planetary path.

An s sub-shell contains one orbital with a spherical shape. The size of the s orbital increases with principal shell number, but its shape remains spherical in this model.

A p sub-shell contains three orbitals, labelled pₓ, pᵧ and p_z. Each has a dumbbell-like shape and is oriented along one of three mutually perpendicular axes; in a ground-state atom the three p orbitals in one sub-shell have equal energy.

Use the shapes to describe allowed regions and spatial orientation, not fixed electron tracks. The label 2p identifies shell n = 2 and the p sub-shell; it does not change the orbital shapes into separate shells.

At this level, recognise and compare s and p shapes without adding detailed d-orbital drawings. Shape, orbital count and electron occupancy are related but distinct properties.

A free radical is a species containing one or more unpaired electrons. The defining evidence is the electron arrangement, not whether the species is neutral or charged.

In an electron-box diagram, an unpaired electron is shown by a single arrow in an orbital; in a configuration, inspect the outer sub-shell occupancy. A dot can also represent the unpaired electron in a radical symbol.

A chlorine radical has a 3p⁵ outer arrangement: two 3p orbitals contain paired electrons and one 3p orbital contains one unpaired electron. The single electron is the evidence for radical status.

Do not equate radical with ion: a neutral species can be a radical, and charge alone does not establish an unpaired electron. The SME also links radical formation to homolytic fission, where a covalent-bond electron pair separates evenly; do not extend this card beyond that supported boundary.