Structure of the Atom and the Periodic Table
What the Bagrut tests here
Atomic structure questions ask you to move between three representations: the electron configuration, the element's position in the periodic table, and physical trends such as atomic radius, first ionization energy, electronegativity and electron affinity. You are usually given a proton number and asked to explain a trend, not just state it.
Core rules
- Aufbau order: 1s 2s 2p 3s 3p 4s 3d 4p. The 4s orbital fills before 3d.
- Pauli: two electrons per orbital, opposite spins. Hund: within a degenerate set (three p or five d orbitals), occupy singly with parallel spins before pairing.
- Two configurations the exam likes: Cr is [Ar]3d5 4s1 and Cu is [Ar]3d10 4s1, because half-filled and filled d subshells are stabilized.
- Across a period, effective nuclear charge rises, so radius falls while ionization energy and electronegativity rise.
- Down a group, n increases, so radius rises, ionization energy falls and metallic character rises.
- Isoelectronic species: the one with more protons is smaller, so Na+ < Ne < F-.
Worked example
Why is the first ionization energy of Mg (738 kJ/mol) higher than that of Al (578 kJ/mol), even though Al has the larger nuclear charge?
Mg loses a 3s electron from a filled 3s2 subshell. Al loses a 3p electron, which lies higher in energy and is shielded by the 3s2 pair, so it is easier to remove. The trend across a period is therefore not monotonic: the group 2 to group 13 drop is a standard exception, as is the group 15 to group 16 drop caused by pairing an electron in a p orbital (N > O).
Traps that cost marks
Stating that ionization energy always increases across a period without the Mg/Al and N/O exceptions. Writing 3d before 4s in a configuration. Confusing electron affinity with electronegativity. Comparing radii without checking that the species are isoelectronic or in the same group.