What you'll learn
Electronic configuration describes how the electrons in an atom are arranged in energy levels, sub-levels and orbitals, and periodicity is the regular repetition of properties across the Periodic Table that this arrangement produces. This guide covers the evidence for energy levels from atomic emission spectra, the shapes of s and p orbitals, the aufbau principle, Hund's rule and the Pauli exclusion principle, writing configurations for atoms and ions (including the exceptions chromium and copper), ionisation energy and the factors that affect it, and trends in atomic radius, ionisation energy and electronegativity across periods and down groups. Understanding these ideas explains why elements behave as they do, from the reactivity of sodium to the inertness of argon.
CAPE questions frequently ask you to interpret graphs of successive ionisation energies or first ionisation energies across a period, and to explain trends using nuclear charge, shielding and distance. Precise explanations using these three factors earn full marks.
Key terms and definitions
Energy level (shell) — A region around the nucleus where electrons have a particular energy, labelled by principal quantum number n = 1, 2, 3…
Sub-level — A division of an energy level: s, p, d or f.
Orbital — A region of space that can hold up to two electrons with opposite spins.
Aufbau principle — Electrons fill orbitals in order of increasing energy.
Hund's rule — Electrons occupy orbitals of equal energy singly, with parallel spins, before pairing.
Pauli exclusion principle — No orbital can hold more than two electrons, and those two must have opposite spins.
First ionisation energy — The energy required to remove one mole of electrons from one mole of gaseous atoms to form one mole of gaseous 1+ ions: X(g) → X⁺(g) + e⁻.
Successive ionisation energies — The energies needed to remove each further electron in turn.
Shielding — The reduction in attraction between the nucleus and outer electrons caused by inner electrons.
Effective nuclear charge — The net positive charge experienced by an outer electron after shielding.
Atomic radius — Half the distance between the nuclei of two bonded atoms of the same element.
Electronegativity — The ability of an atom in a covalent bond to attract the bonding pair of electrons.
Core concepts
Evidence for energy levels: emission spectra
When atoms are heated or excited by electricity, electrons absorb energy and move to higher energy levels. When they fall back, they emit energy as light of specific frequencies, producing a line emission spectrum rather than a continuous rainbow. Each line corresponds to a transition between two levels, with energy ΔE = hf, where h is Planck's constant and f is the frequency.
In the hydrogen spectrum:
- transitions to n = 1 produce the Lyman series (ultraviolet);
- transitions to n = 2 produce the Balmer series (visible);
- lines converge at higher frequency because energy levels get closer together further from the nucleus.
The convergence limit of the Lyman series corresponds to the electron being removed completely, so it can be used to calculate hydrogen's ionisation energy. Line spectra are evidence that electrons can only occupy fixed (quantised) energy levels.
Orbitals and sub-levels
| Sub-level | Number of orbitals | Maximum electrons |
|---|---|---|
| s | 1 | 2 |
| p | 3 | 6 |
| d | 5 | 10 |
| f | 7 | 14 |
An s orbital is spherical; a p orbital is dumb-bell shaped, with three p orbitals (pₓ, p_y, p_z) at right angles to each other.
Energy level n can hold 2n² electrons: 2 in n = 1, 8 in n = 2, 18 in n = 3.
Filling order and configurations
The order of filling is: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p… The 4s sub-level fills before 3d because it is slightly lower in energy when empty.
Examples:
- Na: 1s² 2s² 2p⁶ 3s¹
- Cl: 1s² 2s² 2p⁶ 3s² 3p⁵
- Fe: 1s² 2s² 2p⁶ 3s² 3p⁶ 3d⁶ 4s²
Exceptions: chromium is [Ar] 3d⁵ 4s¹ and copper is [Ar] 3d¹⁰ 4s¹, because a half-filled or completely filled d sub-level gives extra stability.
Ions: for transition metals, the 4s electrons are removed before the 3d electrons when forming positive ions. So Fe²⁺ is [Ar] 3d⁶ and Fe³⁺ is [Ar] 3d⁵.
Orbital diagrams (boxes with arrows) show Hund's rule: nitrogen's three 2p electrons occupy separate orbitals with parallel spins.
Factors affecting ionisation energy
Three factors determine how strongly an outer electron is held:
- Nuclear charge — more protons means stronger attraction and higher ionisation energy.
- Distance from the nucleus (atomic radius) — electrons further away are held less strongly.
- Shielding — more inner electrons reduce the attraction felt by outer electrons.
A fourth consideration is electron pairing: electrons in a doubly occupied orbital repel each other, making one easier to remove.
Trend in first ionisation energy across Period 3
First ionisation energy generally increases from sodium to argon because nuclear charge increases while shielding stays roughly the same (electrons are added to the same shell), so the outer electrons are held more strongly and the atomic radius decreases.
There are two dips:
- Mg → Al: aluminium's outer electron is in a 3p orbital, which is higher in energy and slightly further from the nucleus than magnesium's 3s, and is partly shielded by the 3s electrons, so it is easier to remove.
- P → S: in sulfur, the fourth 3p electron must pair with another in the same orbital. Repulsion between the paired electrons makes one easier to remove than in phosphorus, where each 3p orbital holds one electron.
Trend down a group
First ionisation energy decreases down a group. Although nuclear charge increases, the outer electrons are in shells further from the nucleus and are more shielded by inner electrons, so they are held less strongly. This explains why reactivity of Group I metals increases down the group.
Successive ionisation energies
Successive ionisation energies always increase, because each electron is removed from an increasingly positive ion. A large jump occurs when an electron is removed from a new shell closer to the nucleus. For aluminium, the first three ionisation energies rise steadily, then the fourth is much larger, showing that aluminium has three outer electrons and is in Group III.
Plotting the logarithm of successive ionisation energies against electron number reveals the shell structure clearly, with jumps between shells.
Atomic radius and electronegativity
Across a period: atomic radius decreases because increasing nuclear charge pulls the electrons in the same shell closer; electronegativity increases for the same reason.
Down a group: atomic radius increases because extra shells are added; electronegativity decreases because the bonding electrons are further from the nucleus and more shielded.
Fluorine is the most electronegative element. Electronegativity differences between bonded atoms determine bond polarity, linking this topic to bonding.
Worked examples
Example 1: Writing configurations
Write the full electronic configurations of S and S²⁻, and of Cu and Cu²⁺.
S: 1s² 2s² 2p⁶ 3s² 3p⁴. S²⁻ gains two electrons: 1s² 2s² 2p⁶ 3s² 3p⁶. Cu: 1s² 2s² 2p⁶ 3s² 3p⁶ 3d¹⁰ 4s¹. Cu²⁺ loses the 4s electron then one 3d electron: 1s² 2s² 2p⁶ 3s² 3p⁶ 3d⁹.
Example 2: Identifying a group from ionisation energies
An element has successive ionisation energies (kJ mol⁻¹): 738, 1451, 7733, 10 540. Which group is it in?
The large jump occurs between the second (1451) and third (7733) ionisation energies, so the element has two outer electrons and is in Group II (it is magnesium).
Example 3: Explaining a dip
Explain why the first ionisation energy of sulfur is lower than that of phosphorus. (3 marks)
Phosphorus has the configuration 3p³ with one electron in each 3p orbital (1). In sulfur (3p⁴), two electrons share one 3p orbital (1). Repulsion between the paired electrons makes it easier to remove one of them, so less energy is needed (1).
Example 4: Emission spectra
Explain why the lines in the hydrogen emission spectrum converge at higher frequencies.
The energy levels in the atom become closer together as n increases. Transitions from higher levels to the same lower level therefore differ by smaller amounts of energy, so the lines get closer together and eventually converge.
Example 5: Ionisation energy from the hydrogen spectrum
The convergence limit of the Lyman series in the hydrogen spectrum is at a frequency of 3.27 × 10¹⁵ Hz. Calculate the first ionisation energy of hydrogen in kJ mol⁻¹. (h = 6.63 × 10⁻³⁴ J s; L = 6.02 × 10²³ mol⁻¹)
Energy per atom: E = hf = 6.63 × 10⁻³⁴ × 3.27 × 10¹⁵ = 2.17 × 10⁻¹⁸ J. Energy per mole: 2.17 × 10⁻¹⁸ × 6.02 × 10²³ = 1.31 × 10⁶ J mol⁻¹ = 1310 kJ mol⁻¹ (the accepted value is 1312 kJ mol⁻¹).
The convergence limit represents the electron being promoted from n = 1 to n = ∞ — that is, removed completely.
Example 6: Stability of Fe³⁺ compared with Fe²⁺
Use electronic configurations to explain why iron(II) compounds are readily oxidised to iron(III) in air.
Fe²⁺ is [Ar] 3d⁶ and Fe³⁺ is [Ar] 3d⁵. Removing one more electron gives Fe³⁺ a half-filled d sub-level, with one electron in each of the five 3d orbitals, which is a particularly stable arrangement. This contributes to the ease with which Fe²⁺ is oxidised to Fe³⁺ — for example, the green iron(II) hydroxide precipitate turns brown on standing in air.
Example 7: Interpreting a graph of first ionisation energies
On a graph of first ionisation energy against atomic number for elements 1–20, sharp drops occur after helium, neon and argon. Explain these drops.
Each drop occurs at the start of a new period (Li, Na, K). The outer electron is in a new, higher shell, further from the nucleus and more shielded by inner electrons, so it is held much less strongly and less energy is needed to remove it, despite the increase in nuclear charge.
Example 8: Ionic radii in an isoelectronic series
Arrange N³⁻, O²⁻, F⁻, Na⁺ and Mg²⁺ in order of decreasing radius and explain.
All five ions have the same electron configuration, 1s² 2s² 2p⁶ (they are isoelectronic). The number of protons increases from 7 (N) to 12 (Mg), so the nuclear attraction for the same number of electrons increases and the ions get smaller: N³⁻ > O²⁻ > F⁻ > Na⁺ > Mg²⁺.
Periodic trends summary
| Property | Across a period (left → right) | Down a group |
|---|---|---|
| Atomic radius | decreases | increases |
| First ionisation energy | generally increases (dips at Groups III and VI) | decreases |
| Electronegativity | increases | decreases |
| Metallic character | decreases | increases |
All these trends can be explained using nuclear charge, distance from the nucleus and shielding.
Common mistakes and how to avoid them
- Filling 3d before 4s when writing atomic configurations, or removing 3d before 4s when forming ions.
- Forgetting the chromium and copper exceptions.
- Explaining trends with only one factor. Use nuclear charge, distance and shielding together.
- Saying ionisation energy decreases across a period because of shielding. Shielding is roughly constant across a period; it is nuclear charge that increases.
- Omitting state symbols in ionisation energy equations — (g) is essential.
How this links to your Internal Assessment
Although electronic structure is mainly theoretical, it supports practical work on flame tests, where metal ions produce characteristic colours as excited electrons fall to lower energy levels: sodium gives yellow, potassium lilac, calcium brick-red and copper blue-green. In an SBA on identifying ions, you can explain your observations using emission of specific frequencies of light, linking your practical evidence to atomic theory.
Exam technique for Electronic configuration and periodicity
- For ionisation energy explanations, structure answers using nuclear charge, distance and shielding, and mention pairing or sub-level for the dips.
- Write equations for ionisation energy with state symbols.
- For graphs of successive ionisation energies, identify the big jumps and relate them to shells and group.
- Use noble gas shorthand ([Ar]) only when the question allows it.
Quick revision summary
- Line emission spectra show electrons occupy quantised energy levels; Lyman (UV) and Balmer (visible) series; convergence limit gives ionisation energy.
- s (spherical, 2 electrons), p (dumb-bell, 6), d (10), f (14); filling order 1s 2s 2p 3s 3p 4s 3d 4p.
- Cr = [Ar] 3d⁵ 4s¹; Cu = [Ar] 3d¹⁰ 4s¹; 4s electrons are removed first in ions.
- Ionisation energy depends on nuclear charge, distance, shielding and electron pairing.
- Across a period: IE increases (dips at Al and S), radius decreases, electronegativity increases.
- Down a group: IE decreases, radius increases, electronegativity decreases.
- Big jumps in successive IEs reveal the number of outer electrons and the group.