Part 4: Ionisation energies and experimental evidence

Use precise gaseous equations and linked explanations to infer shells, subshells and group membership from energy data.

Ionisation energy specifies both the species and its state

First ionisation energy is the energy required to remove one electron from each atom in one mole of gaseous atoms, forming one mole of gaseous 1+ ions. It is an endothermic process, quoted in kJ mol⁻¹. Gaseous atoms are specified so the value measures electron removal without also melting a solid or breaking molecular bonds.

The second ionisation energy removes one electron from each ion in one mole of gaseous 1+ ions to form one mole of gaseous 2+ ions. More generally, the nth ionisation removes an electron from X⁽ⁿ⁻¹⁾⁺(g), not from a fresh neutral atom. The equation for second ionisation must not remove two electrons from the neutral atom.

Name the electron and the competing attractions

A greater nuclear charge tends to attract an electron more strongly. A greater distance from the nucleus and greater shielding by other electrons tend to reduce that attraction. An electron in a higher-energy subshell is generally easier to remove. Explain a comparison using the electron actually removed, rather than saying only that one atom is more stable.

Across a main-group period, proton number rises while added electrons enter the same principal shell. Shielding by the inner shells is similar, so attraction increases, atomic radius generally decreases and first ionisation energy generally rises. Do not claim shielding is exactly zero or every adjacent step rises.

Down a group, the outer electron occupies a shell farther from the nucleus and is more shielded by inner electrons. These effects outweigh the greater nuclear charge, so first ionisation energy falls. For Group 2, compare removal from 3s in magnesium with 4s in calcium: calcium’s outer electron is farther away and more shielded.

A large jump identifies the change to an inner shell

Successive ionisation energies for the same element increase: fewer electrons remain around the same nucleus and the ion becomes more positive. A particularly large increase occurs when the next electron must come from a lower principal shell, closer to the nucleus and less shielded. Distinguish this shell change from the ordinary increases within one outer shell.

Constructed data in kJ mol⁻¹ are 600, 1200, 7500, 9800 and 12 000. The large increase is between the second and third removals, so there were two outer electrons: a main-group Group 2 element. It is the number removed before the jump that gives the outer-electron count; the third value does not mean Group 3.

On a logarithmic plot, multiplicative jumps remain clear when values span a wide range. Here log₁₀(600) = 2.778, log₁₀(1200) = 3.079 and log₁₀(7500) = 3.875. The second increase is 0.796 logarithmic units, representing a factor 6.25, versus a factor two for the first. Always check whether an axis is linear or logarithmic.

Small dips reveal subshell energy and pairing

Boron’s first ionisation energy is lower than beryllium’s. Be loses a 2s electron, whereas B loses a higher-energy 2p electron, which is somewhat shielded by 2s. This change outweighs the increase in proton number. The analogous Period 3 dip is Mg to Al, involving 3s and 3p.

Oxygen’s first ionisation energy is lower than nitrogen’s. Nitrogen has three singly occupied 2p orbitals; oxygen has one paired 2p orbital. Repulsion between the two electrons in that orbital makes one easier to remove. Both elements lose 2p electrons, so a change from s to p is not the explanation. The analogous Period 3 dip is P to S.

The two dips are separate pieces of evidence: subshell energy explains Be/B and Mg/Al; electron pairing explains N/O and P/S. Use the correct principal-shell number when transferring the explanation to another period.

Line spectra provide independent evidence for discrete energies

Energy supplied by a flame or electrical discharge can excite electrons to higher allowed energy levels. When an excited electron falls to a lower level, a photon is emitted with energy equal to the gap between those levels. Only particular gaps are allowed, so the spectrum contains separate lines rather than every wavelength.

This supports quantised electronic energy levels. Successive-ionisation jumps provide complementary evidence for shells, and the smaller first-ionisation anomalies support subshells and electron pairing. A model is supported by its ability to explain several observations; electrons do not need to be pictured as miniature planets on fixed tracks.

With a hand-held spectroscope, view the light from a supervised flame test without looking directly at an intense source. Compare the positions of distinct lines, not only the overall flame colour. Repeat with clean apparatus because sodium contamination can dominate the visible colour. The detailed explanation of flame tests is developed in Topic 4.

Build a comparison with a complete causal chain

For an ‘explain’ question, connect the change in electronic environment to attraction and then energy. For calcium versus magnesium: an extra occupied shell places the removed electron farther from the nucleus and adds shielding; despite more protons, attraction to the outer electron is weaker, so less energy is needed.

Pearson 8CH0/01 June 2023 Q9(b)(ii) asked for both the down-group trend and the comparison between first and second ionisation energies. Its scheme assesses linked reasoning as well as chemical content. The report notes that some answers covered only the down-group trend. Read every comparison requested; six unrelated phrases do not automatically earn six marks.

Quick checks

Original Finesse questions. Reveal the indicative worked solutions after attempting each question; these are not official Edexcel mark allocations.

Sources