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1.3 Electrons, energy levels and atomic orbitals

Syllabus
9701–2028–2029
Topic
1.3
Level
AS

Learning objectives

1.3.0Scope note• In 1.3 each atom or ion described will be in the ground state. Only the elements hydrogen to krypton will be assessed.1.3.1Terms• Know terms:: shells, sub-shells and orbitals; principal quantum number (n); ground state, limited to electronic configuration1.3.2Number of orbitals making up s, p and d• Describe the number of orbitals making up s, p and d sub-shells, and the number of electrons that can fill s, p and d sub-shells1.3.3Order of increasing energy of the sub-shells• Describe the order of increasing energy of the sub-shells within the first three shells and the 4s and 4p sub-shells1.3.4Electronic configurations to include• Describe the electronic configurations to include the number of electrons in each shell, sub-shell and orbital1.3.5Electronic configurations• Explain the electronic configurations in terms of energy of the electrons and inter-electron repulsion1.3.6Determine the electronic configuration• Determine the electronic configuration of atoms and ions given the atomic/proton number and charge, using either of the following conventions: e.g. for Fe: 1s22s22p63s23p63d64s2 (full electronic configuration) or [Ar] 3d64s2 (shorthand electronic configuration)1.3.7The electrons in boxes notation• Understand and use the electrons in boxes notation e.g. for Fe: [Ar]1.3.8And sketch the shapes of s and p orbitals• Describe and sketch the shapes of s and p orbitals1.3.9Free radical as a species with one or more• Describe a free radical as a species with one or more unpaired electrons

The assessed 1.3 cases are ground-state species from H to Kr

Throughout section 1.3, every atom or ion is in its ground state: its electrons occupy the lowest-energy arrangement allowed. The assessed elements run from hydrogen, H, to krypton, Kr.

Use this boundary when choosing examples and writing configurations. Excited-state arrangements and elements beyond krypton are outside this section; the boundary limits what is assessed but does not redefine shells, sub-shells or orbitals.

Shells, sub-shells and orbitals form a hierarchy

A shell is a principal energy level identified by the principal quantum number, n. Each shell contains one or more sub-shells, labelled s, p or d within the assessed range, and each sub-shell contains one or more orbitals.

An orbital is a region of space that can hold a maximum of two electrons. It is not a fixed path followed by an electron. For example, 2p identifies the p sub-shell in the shell with n = 2.

The ground state is the lowest-energy electronic configuration available to the species. Do not use shell, sub-shell and orbital as interchangeable terms: they are nested levels of the model.

Orbital count determines each sub-shell's electron capacity

Each orbital holds at most two electrons, so a sub-shell's capacity is twice its number of orbitals.

Sub-shell Number of orbitals Maximum electrons
s 1 2
p 3 6
d 5 10

Keep the two counts distinct: a p sub-shell contains three orbitals but can contain six electrons. This objective assesses s, p and d; adding f is not needed for the stated learning outcome.

Use the assessed sub-shell energy order

A ground-state configuration places electrons into lower-energy sub-shells before higher-energy sub-shells. Sub-shell energies overlap, so principal shell number alone does not give the filling order.

1s<2s<2p<3s<3p<4s<3d<4p1s < 2s < 2p < 3s < 3p < 4s < 3d < 4p

Potassium therefore ends in 4s¹, not 3d¹. The sequence is an energy order for the sub-shells required here; it is not the simple numerical order 1, 2, 3, 4.

A configuration records electrons at shell, sub-shell and orbital levels

In a term such as 2p⁴, 2 is the principal shell, p is the sub-shell, and the superscript 4 is the number of electrons in that sub-shell. The superscripts in a complete configuration must add to the species' total number of electrons.

O: 1s2 2s2 2p4\mathrm{O}:\ 1s^2\,2s^2\,2p^4

Sub-shell notation gives the total within each sub-shell. Electron-in-box notation adds the orbital-level detail by showing how those electrons are distributed among the individual orbitals.

Do not read the superscript as an orbital count or an atomic number. It is an electron count for that sub-shell.

Lowest energy and minimum repulsion explain electron placement

In the ground state, electrons occupy the available sub-shells in increasing energy. Within one sub-shell, orbitals have the same energy, so electrons occupy separate orbitals before pairing.

Separate occupancy keeps electrons farther apart and reduces inter-electron repulsion. Once every orbital in that sub-shell contains one electron, further electrons must pair; the two electrons in one orbital have opposite spins.

For 2p³, place one electron in each of the three 2p orbitals. For 2p⁴, the fourth electron pairs in one orbital. Pairing earlier would give greater repulsion without lowering the sub-shell energy.

Determine configurations by count, fill, then check

First determine the electron count: a neutral atom has Z electrons; subtract the positive charge for a cation or add the magnitude of the negative charge for an anion. Fill sub-shells in the assessed energy order, then check that all superscripts add to this count.

Fe: 1s2 2s2 2p6 3s2 3p6 3d6 4s2=[Ar] 3d6 4s2\mathrm{Fe}:\ 1s^2\,2s^2\,2p^6\,3s^2\,3p^6\,3d^6\,4s^2 = [Ar]\,3d^6\,4s^2

Fe2+: [Ar] 3d6\mathrm{Fe^{2+}}:\ [Ar]\,3d^6

When a transition-metal ion forms, remove 4s electrons before 3d electrons. Changing charge changes the electron count, not the proton number or the identity of the element.

Electron-in-box notation shows occupancy and spin

Each box represents one orbital and each arrow represents one electron. Two arrows in one box must point in opposite directions; among equal-energy orbitals, place one electron in each box before pairing.

2p orbital first second third
2p⁴ occupancy ↑↓ ↑ ↑

This 2p⁴ diagram contains four electrons: one pair and two unpaired electrons. A box is an orbital, not a shell, and arrow direction represents spin rather than electron motion along a path.

s orbitals are spherical; p orbitals are directional

An s orbital is spherical around the nucleus. A larger principal quantum number gives a larger s orbital, while the assessed overall shape remains spherical.

Each p sub-shell contains three dumbbell-shaped orbitals, labelled pₓ, pᵧ and p_z. Their lobes point along three mutually perpendicular axes, and the nucleus lies at the centre between the two lobes of each orbital.

In a sketch, show one sphere for an s orbital or two equal lobes on a straight axis for a p orbital, centred on the nucleus. The boundary is a probability-region model, not a hard surface or an electron track.

A free radical has at least one unpaired electron

A free radical is a species with one or more unpaired electrons. Radical status is determined by electron pairing, not by whether the species is neutral or charged.

Cl⋅\mathrm{Cl\boldsymbol{\cdot}}

The dot represents the unpaired electron. In electron-in-box notation, the same evidence is a singly occupied orbital. A lone pair contains two paired electrons, so a lone pair alone does not make a species a radical.