• Emission spectra are produced when photons are emitted from atoms as excited electrons return to a lower energy level.

• The line emission spectrum of hydrogen provides evidence for the existence of electrons in discrete energy levels, which converge at higher energies.

• The main energy level or shell is given an integer number, n, and can hold a maximum number of electrons, 2n².

• A more detailed model of the atom describes the division of the main energy level into s, p, d and f sub-levels of successively higher energies.

• Sub-levels contain a fixed number of orbitals, regions of space where there is a high probability of finding an electron.

• Each orbital has a defined energy state for a given electronic configuration and chemical environment and can hold two electrons of opposite spin.

• Description of the relationship between colour, wavelength, frequency and energy across the electromagnetic spectrum.

• Distinction between a continuous spectrum and a line spectrum.

• Description of the emission spectrum of the hydrogen atom, including the relationships between the lines and energy transitions to the first, second and third energy levels.

• Recognition of the shape of an s atomic orbital and the px, py and pz atomic orbitals.

• Application of the Aufbau principle, Hund’s rule and the Pauli exclusion principle to write electron configurations for atoms and ions up to Z = 36.

• Details of the electromagnetic spectrum are given in the data booklet in section 3.

• The names of the different series in the hydrogen line emission spectrum are not required.

• Full electron configurations (eg 1s22s22p63s23p4) and condensed electron configurations (eg [Ne] 3s23p4) should be covered.

• Orbital diagrams should be used to represent the character and relative energy of orbitals. Orbital diagrams refer to arrow-in-box diagrams.

• The electron configurations of Cr and Cu as exceptions should be covered.