Electron configuration — AQA A-Level Chemistry
Test yourself on Electron configuration with AQA A-Level practice questions.
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Electron configuration explained
Electrons fill sub-shells in order of increasing energy: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p.
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An s sub-shell holds 2 electrons, p holds 6 and d holds 10, because each orbital takes two electrons of opposite spin. The 4s sub-shell lies just below 3d in energy so it fills first, but once the 3d sub-shell is occupied it drops below 4s, which is why transition metal ions lose their 4s electrons first: iron is 1s² 2s² 2p⁶ 3s² 3p⁶ 3d⁶ 4s² but iron(II) is 3d⁶ with no 4s. Chromium and copper break the pattern, adopting 3d⁵ 4s¹ and 3d¹⁰ 4s¹ because a half-filled or full d sub-shell is more stable.
Ionisation energies.
The first ionisation energy is the energy needed to remove one electron from each atom in one mole of gaseous atoms, forming one mole of gaseous 1⁺ ions. It is always endothermic because work is done against the electrostatic attraction between the nucleus and the electron. Three factors control its size: nuclear charge, which increases attraction; atomic radius, since a more distant electron is held less strongly; and shielding by inner shells, which reduces the effective pull. Successive ionisation energies always rise, partly because each electron after the first is pulled away from an ion that is already positive and smaller, so the remaining electrons are held more tightly.
Students should be able to: define first ionisation energy write equations for first and successive ionisation energies explain how first and successive ionisation energies in Period 3 (Na–Ar) and in Group 2 (Be–Ba) give evidence for electron configuration in sub-shells and in shells.
First ionisation energy is the energy required to remove one mole of electrons from one mole of gaseous atoms to form one mole of gaseous 1+ ions. Write equations with state symbols, e.g. Na(g) → Na⁺(g) + e⁻. Across Period 3, values rise overall as nuclear charge increases while shielding stays similar, but there are two dips. Aluminium falls below magnesium because its outer electron is in a 3p sub-shell, higher in energy and shielded by the 3s pair. Sulfur falls below phosphorus because its 3p⁴ arrangement forces two electrons into one orbital, causing mutual repulsion. Down Group 2, values fall as radius and shielding increase. A large jump between successive values marks a change to a shell nearer the nucleus.
Your focus
- Write the full sub-shell electron configuration of any atom or ion up to krypton, adjusting the electron count for charge.
- Explain why iron(II) is written with no 4s electrons despite 4s filling before 3d in the atom.
- Complete an electrons-in-boxes diagram showing single occupancy before pairing and opposite spins in a filled orbital.
Show all 9 objectives
- Define first ionisation energy in a form that specifies one mole, the gaseous state and the charge on the product ion.
- Explain, using nuclear charge, radius and shielding, why one element has a higher first ionisation energy than another.
- Justify why ionisation is always an endothermic change.
- Define first ionisation energy and write equations, with state symbols, for first and successive ionisation energies.
- Deduce the group of an element from the position of the large jump in its successive ionisation energies.
- Explain the dips at aluminium and at sulfur in terms of sub-shell energy and electron pairing respectively.
Electron configuration exam tips
Marking Points
- one mark for the sub-shells written in the correct filling order including 4s before 3d
- one mark for the correct number of electrons in each sub-shell totalling the atomic number
- one mark for removing 4s electrons before 3d electrons when writing a transition metal ion
- one mark for the exceptional configurations of chromium and copper
- one mark for removal of one electron from each atom in one mole of gaseous atoms
- one mark for the product described as one mole of gaseous 1⁺ ions
- one mark for naming nuclear charge, atomic radius and shielding as the controlling factors
- one mark for explaining a rise in successive values by removal from an increasingly positive ion
- one mark for defining first ionisation energy as the energy required to remove one mole of electrons from one mole of gaseous atoms
- one mark for an equation showing one electron removed with gaseous state symbols on both species
- one mark for the large jump between successive values identifying a change of shell
- one mark for the dip at aluminium explained by the outer electron occupying a 3p sub-shell of higher energy
- one mark for the dip at sulfur explained by repulsion between the paired electrons in one 3p orbital
Examiner Tips
- 💡Count the electrons in your finished configuration and check the total against the atomic number or adjusted ion charge.
- 💡Write the atom first, then edit it to make the ion; do not build an ion configuration from scratch.
- 💡In electrons-in-boxes diagrams, fill each orbital singly before pairing, and show paired arrows pointing opposite ways.
- 💡Learn the definition as a single sentence containing the words one mole, gaseous and 1⁺ ions.
- 💡Always add state symbols to ionisation equations; they carry a mark of their own.
- 💡When comparing two values, mention all three factors and say which one dominates.
- 💡Count the jumps, not the values: the number of electrons before the first big jump gives the group number, except for Group 0 where it is 8.
- 💡Name the sub-shell, 3s or 3p, whenever you explain a Period 3 dip; a vague reference to energy levels will not score.
- 💡For the second ionisation energy always start the equation from the 1⁺ ion, never from the atom.
Common Mistakes
- removing 3d electrons first when forming an ion such as iron(II)
- writing chromium as 3d⁴ 4s² instead of 3d⁵ 4s¹
- putting more than six electrons in a p sub-shell or more than ten in a d sub-shell
- inventing a 2d sub-shell, which does not exist
- adding electrons to a positive ion rather than removing them
- omitting the gaseous state from the definition or from the equation
- defining it as the energy to remove all the electrons rather than one per atom
- describing ionisation as exothermic because an ion is formed
- referring to one atom rather than one mole so the quantity has no molar units
- writing the equation without gaseous state symbols
- showing two electrons removed in a first ionisation energy equation
- explaining the aluminium dip by extra shielding from an additional shell rather than the 3s pair and higher 3p energy
- attributing the sulfur dip to a change in nuclear charge instead of paired-electron repulsion