CIE A Level Chemistry Atomic Structure Exam Questions

About CIE A Level Chemistry Atomic Structure

Atomic structure is the first topic in the Cambridge International AS & A Level Chemistry (9701) syllabus and underpins almost everything that follows. You need to describe protons, neutrons and electrons, explain how they behave in an electric field, and write electronic configurations using shells, subshells and orbitals. The topic then moves on to ionisation energies, which give evidence for this electron structure and explain trends across periods and down groups. It is examined as AS content, and it links directly to chemical bonding, periodicity and the transition elements. Examiners regularly ask for full configurations of ions, explanations of ionisation energy trends and interpretation of successive ionisation energy data.

What you need to know

• Particles and isotopes Protons (relative charge +1, mass 1), neutrons (0, mass 1) and electrons (-1, mass 1/1836). Isotopes have the same number of protons but different numbers of neutrons, so they share chemical properties but differ in physical properties such as density.
• Behaviour in an electric field Protons are deflected towards the negative plate, electrons towards the positive plate, and neutrons are undeflected. Electrons are deflected more because the angle of deflection depends on charge/mass ratio.
• Shells, subshells and orbitals Each orbital holds up to two electrons with opposite spins. s orbitals are spherical and p orbitals are dumbbell shaped. Subshells fill in the order 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, and the 4s electrons are lost before 3d electrons when transition metal ions form.
• Exceptions in configurations Chromium is [Ar]3d5 4s1 and copper is [Ar]3d10 4s1, because a half-filled or full 3d subshell is more stable. You should also be able to use box notation and the shorthand [Ne] or [Ar].
• First ionisation energy The energy required to remove one electron from each atom in one mole of gaseous atoms to form one mole of gaseous 1+ ions, for example Na(g) → Na+(g) + e-. It depends on nuclear charge, distance of the electron from the nucleus and shielding by inner electrons.
• Trends and evidence First ionisation energy generally increases across a period and decreases down a group. Dips at Al (electron in a higher-energy 3p subshell) and S (spin-pair repulsion in a 3p orbital) are evidence for subshells. Large jumps in successive ionisation energies show where a new, inner shell begins.
• Atomic and ionic radius Atomic radius decreases across a period as nuclear charge increases with similar shielding. Cations are smaller than their atoms and anions are larger, because of changes in electron number and repulsion.

Common mistakes in Atomic Structure

• Writing ionisation equations without state symbols; the atom and ion must both be shown as (g).
• Removing 3d electrons before 4s when forming transition metal ions; Fe2+ is [Ar]3d6, not [Ar]3d4 4s2.
• Explaining the dip at sulfur by saying the electron is further away; it is due to repulsion between the paired electrons in one 3p orbital.
• Saying neutrons are deflected less than protons; neutrons have no charge and are not deflected at all.
• Stating that ionisation energy increases across a period because of more shells; the number of shells is the same, it is the increasing nuclear charge with similar shielding that matters.

Try three Atomic Structure exam-style questions