Describe electron configuration in terms of shells, subshells (s, p, d) and orbitals using the (1s 2s 2p 3s 3p 4s 3d 4p) filling order, and explain the periodic trends in atomic radius, first ionisation energy and electronegativity using effective nuclear charge and shielding

What this dot point is asking

QCAA wants you to write electron configurations using s, p and d subshells with the standard filling order, predict the configuration of ions (including the 4s-removed-first rule for transition metals), and explain the three core periodic trends (atomic radius, first ionisation energy, electronegativity) in terms of effective nuclear charge and shielding. This is the foundation for every bonding argument in Topic 2.

The answer

Electrons occupy shells (energy levels labelled by principal quantum number n) made up of subshells (s, p, d, f) that contain orbitals. Orbitals fill in a specific order set by their energies, following the aufbau principle. Periodic trends arise because of how the effective nuclear charge experienced by the outer electrons changes across periods and down groups.

Shells, subshells and orbitals

The capacity of shell n is $2n^2$ electrons. Shell 1: 2 (1s). Shell 2: 8 (2s, 2p). Shell 3: 18 (3s, 3p, 3d). Shell 4: 32 (4s, 4p, 4d, 4f).

Each orbital holds up to two electrons of opposite spin (Pauli exclusion). Within a subshell, orbitals are filled singly before pairing (Hund's rule), which minimises electron-electron repulsion.

The filling order

QCE Chemistry expects you to use the standard aufbau order. For Z up to 36 (krypton), the order is:

The 4s subshell fills before 3d in neutral atoms because 4s is at lower energy than 3d at that stage. Beyond Kr the order becomes 5s, 4d, 5p, and so on.

Worked configurations:

• Na (Z = 11): $1s^2 2s^2 2p^6 3s^1$
• Cl (Z = 17): $1s^2 2s^2 2p^6 3s^2 3p^5$
• Ca (Z = 20): $1s^2 2s^2 2p^6 3s^2 3p^6 4s^2$
• Fe (Z = 26): $1s^2 2s^2 2p^6 3s^2 3p^6 4s^2 3d^6$
• Br (Z = 35): IMATH_8

Two QCAA-relevant exceptions (chromium and copper) gain extra stability by half-filled or fully filled d subshells:

• Cr (Z = 24): $1s^2 2s^2 2p^6 3s^2 3p^6 4s^1 3d^5$ (not $4s^2 3d^4$).
• Cu (Z = 29): $1s^2 2s^2 2p^6 3s^2 3p^6 4s^1 3d^{10}$ (not $4s^2 3d^9$).

Electron configurations of ions

For main-group ions, electrons are added or removed from the outermost subshell.

• F (Z = 9): $1s^2 2s^2 2p^5$. F-: $1s^2 2s^2 2p^6$ (noble-gas Ne configuration).
• Mg (Z = 12): $1s^2 2s^2 2p^6 3s^2$. Mg^2+: $1s^2 2s^2 2p^6$.

For transition metals forming cations, remove the s electrons before the d electrons.

• Fe (Z = 26): $...4s^2 3d^6$. Fe^2+: $...3d^6$ (4s lost). Fe^3+: $...3d^5$ (4s and one 3d lost).
• Cu (Z = 29): $...4s^1 3d^{10}$. Cu^+: $...3d^{10}$. Cu^2+: $...3d^9$.

This is a frequent QCAA test point because it contradicts the naive "last in, first out" expectation.

Effective nuclear charge and shielding

Effective nuclear charge (Z_eff) is the net positive charge experienced by an outer electron, after accounting for shielding (also called screening) by inner electrons.

where S is the screening constant. Inner-shell electrons shield outer electrons effectively; electrons in the same subshell shield each other only weakly.

Two principles drive every periodic trend:

1. Across a period: Z rises by one each step; S stays roughly constant (electrons added to the same shell shield each other poorly). So Z_eff rises strongly.
2. Down a group: Z rises but S also rises sharply (a whole new inner shell each step). Z_eff stays roughly constant, but the principal quantum number n rises, so the outermost electrons are physically further from the nucleus.

Atomic radius

Atomic radius is the typical distance from the nucleus to the outermost occupied shell.

• Across a period: radius decreases. Z_eff rises, electrons in the same shell pulled in tighter.
• Down a group: radius increases. Outermost shell has higher n; the electron is intrinsically further out.

Ionic radii follow:

• Cations are smaller than their parent atoms (lose a whole outer shell, or at least reduce electron-electron repulsion).
• Anions are larger than their parent atoms (added electron increases repulsion in the outer shell).

First ionisation energy

First ionisation energy (IE_1) is the energy required to remove the most loosely held electron from a gaseous atom.

• Across a period: IE_1 increases. Higher Z_eff, smaller radius, electron more tightly held.
• Down a group: IE_1 decreases. Outer electron is in a higher shell, further from nucleus, more shielded.

Two subtle "dips" you should be ready to explain:

• Between Group 2 and Group 13 (e.g. Be to B): the outermost electron in B is in a 2p orbital, which is at slightly higher energy than the 2s; easier to remove.
• Between Group 15 and Group 16 (e.g. N to O): N has half-filled 2p (extra stability), so the next electron in O experiences pair-repulsion in the 2p orbital; easier to remove.

These dips are popular QCAA EA short-response prompts.

Electronegativity

Electronegativity is the tendency of an atom in a bond to attract bonding electrons toward itself. Pauling scale; fluorine is set at 4.0, the most electronegative element. Caesium is around 0.7.

• Across a period: electronegativity increases. Same Z_eff reasoning.
• Down a group: electronegativity decreases. Larger atom, bonding electrons further from the nucleus.

The most electronegative elements are concentrated in the upper-right of the periodic table (F, O, N, Cl). The least are bottom-left (Cs, Fr).

Electronegativity governs bond polarity (Topic 2 dot point on covalent bonding) and predicts whether a bond is ionic (large electronegativity difference), polar covalent (moderate difference), or non-polar covalent (small difference).

Putting the three trends together

A single line of reasoning (Z_eff plus shell structure) drives all three. QCAA EA short-response questions often ask you to apply that reasoning to a specific pair, e.g. "Why is the first ionisation energy of Mg higher than that of Na?" or "Why is F more electronegative than Cl?"

Common traps

Filling 3d before 4s in neutral atoms. Aufbau order has 4s first for the neutral atom. Reverse only when forming cations of transition metals.

Forgetting that the s electrons leave first in transition-metal ions. Fe^2+ is 3d^6, not 4s^2 3d^4.

Saying "shielding increases" across a period. Inner-shell shielding is essentially unchanged across a period; only the proton count is changing meaningfully.

Treating the radius and ionisation energy trends as independent. They share a single underlying cause (Z_eff and shell structure). State that cause once and apply it.

Confusing electron affinity with electronegativity. Electron affinity is the energy released when a neutral gas-phase atom gains an electron (a measurable energy). Electronegativity is a dimensionless tendency in a bond. They correlate but are not the same.

In one sentence

Electrons fill subshells in the order 1s 2s 2p 3s 3p 4s 3d 4p obeying aufbau, Pauli and Hund; with s electrons removed before d electrons when forming transition-metal cations, and periodic trends in atomic radius (decreases across, increases down), first ionisation energy (increases across, decreases down) and electronegativity (increases across, decreases down) are all explained by effective nuclear charge rising across periods while shell n rising down groups dominates.