Electron structure — OCR A-Level Chemistry

Electrons in an atom occupy principal energy levels, often called shells, labelled by the principal quantum number n = 1, 2, 3 and so on. Each shell is divided into sub-shells, and the sub-shells available depend on n: shell 1 has only 1s; shell 2 has 2s and 2p; shell 3 has 3s, 3p and 3d; shell 4 has 4s, 4p, 4d and 4f. A sub-shell contains one or more atomic orbitals, each holding up to two electrons of opposite spin. An s sub-shell has one orbital, p has three, d has five and f has seven. Electron configuration fills the lowest available energy sub-shell first, so 4s fills before 3d, giving potassium 1s² 2s² 2p⁶ 3s² 3p⁶ 4s¹.

This criterion asks you to work out how many electrons each of the first four shells can hold. The nth shell contains n² orbitals, and each orbital holds two electrons, so the maximum is 2n². Shell 1 holds 2 electrons, shell 2 holds 8, shell 3 holds 18 and shell 4 holds 32. In practice the third and fourth shells are not always filled to their maximum before the next shell begins, because the 4s sub-shell fills before 3d, but the capacity of each shell is still given by 2n². Learn to calculate 2n² for n = 1 to 4 and to explain the result in terms of sub-shells: shell 1 has 1s; shell 2 has 2s and 2p; shell 3 has 3s, 3p and 3d; shell 4 has 4s, 4p, 4d and 4f.

An atomic orbital is a region around the nucleus where there is a high probability of finding an electron, and it can hold a maximum of two electrons. The statement introduces orbitals as a heading, so the assessed knowledge is the definition itself plus the shapes and labels of the s, p and d orbitals. An s orbital is spherical; the three p orbitals are dumbbell-shaped and mutually perpendicular, labelled px, py and pz; the five d orbitals are more complex. Each orbital holds at most two electrons, and orbitals of equal energy within a subshell are degenerate. A question may ask you to identify a shape, count orbitals in a subshell, or state the maximum electron capacity of a given type of orbital.

This statement defines an atomic orbital precisely: a region around the nucleus where an electron is likely to be found, holding a maximum of two electrons. When two electrons occupy the same orbital they must have opposite spins, shown as one arrow pointing up and one pointing down. This follows the Pauli exclusion principle. In an orbital diagram, a single electron is drawn as an upward arrow; a paired orbital shows an upward and a downward arrow. A question may ask you to complete an orbital diagram, identify a valid spin pairing, or explain why two electrons in one orbital cannot have the same spin.

An orbital is a region of space where the probability of finding an electron is high, and each orbital holds a maximum of two electrons. The s-orbitals are spherical: a 1s orbital is a single sphere, and a 2s orbital is a larger sphere with a spherical node inside it. The p-orbitals are dumbbell-shaped, with two lobes on opposite sides of the nucleus and a nodal plane through the nucleus where electron density is zero. The three p-orbitals of a sub-shell lie at right angles along the x, y and z axes, so they are often drawn as a set of three mutually perpendicular dumbbells. Recognise the shape from a diagram, and distinguish an s-orbital (spherical, one lobe region) from a p-orbital (two lobes, one nodal plane).

Each sub-shell is built from a fixed number of orbitals, and each orbital holds a maximum of two electrons. An s-sub-shell contains one orbital, so it holds a maximum of two electrons. A p-sub-shell contains three orbitals, so it holds a maximum of six electrons. A d-sub-shell contains five orbitals, so it holds a maximum of ten electrons. The pattern is that the number of electrons a sub-shell can hold equals twice the number of its orbitals. Use this relationship to check answers: count the orbitals, then double the number to find the maximum electron capacity. Keep the three sub-shells distinct, since a common slip is to give the orbital count as the electron count.

Filling of orbitals describes the order in which electrons occupy orbitals as the principal quantum number and sub-shell energy rise. Electrons occupy the lowest available energy orbital first, so 1s fills before 2s, and 2s before 2p. Each orbital holds a maximum of two electrons with opposite spins, and within a sub-shell electrons occupy separate orbitals singly before any pairing occurs. For example, the three 2p orbitals each receive one electron before a second electron pairs in any one of them. The 4s orbital fills before 3d because it lies at lower energy in these atoms. A multiple-choice question tests whether you can apply these rules to a given electron configuration or filling sequence rather than recall a single fact.

This criterion asks you to place orbitals in order of increasing energy, covering the first three shells and the 4s and 4p orbitals. The first shell contains only 1s. The second shell contains 2s and 2p. The third shell contains 3s, 3p and 3d. Adding 4s and 4p gives the sequence 1s < 2s < 2p < 3s < 3p < 4s < 3d < 4p, where each symbol is lower in energy than the next. Note that 4s sits below 3d, so the 3d sub-shell is not filled until after 4s. A multiple-choice question tests whether you can order these orbitals correctly and recognise where 4s and 4p fall relative to the third-shell sub-shells.

This rule governs how electrons fill orbitals that have the same energy, such as the three 2p orbitals or the five 3d orbitals. Electrons occupy these degenerate orbitals singly first, each with parallel spins, and only pair up once every orbital of that energy holds one electron. For example, nitrogen's 2p³ arrangement places one electron in each of the three 2p orbitals rather than pairing two in one orbital and leaving another empty. Oxygen adds a fourth 2p electron, which now pairs in one orbital, giving two singly occupied and one doubly occupied 2p orbital. Applying the rule correctly lets you predict unpaired electron counts, which link to magnetic behaviour and to the number of covalent bonds an atom can form.

This outcome asks you to deduce electron configurations rather than simply recall them. You build a configuration by filling subshells in order of increasing energy, respecting the 1s, 2s, 2p, 3s, 3p, 4s, 3d sequence, and by applying the rules that orbitals of equal energy fill singly before pairing and that each orbital holds at most two electrons. You should be able to write configurations in full, such as 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d⁶ for iron, and in shorthand using a noble-gas core, such as [Ar] 4s² 3d⁶. You also deduce configurations of ions by adding or removing electrons, noting that 4s electrons are removed before 3d when transition-metal cations form.

For a neutral atom, the atomic number Z equals the number of protons and therefore the number of electrons. You build the electron configuration by filling sub-shells in order of increasing energy: 1s, then 2s, 2p, 3s, 3p, 4s, then 3d and 4p, which reaches Z = 36 at krypton. Each s sub-shell holds 2 electrons, each p sub-shell holds 6 and each d sub-shell holds 10. For example, Z = 26 (iron) gives 1s² 2s² 2p⁶ 3s² 3p⁶ 3d⁶ 4s², often written as [Ar] 3d⁶ 4s². Remember that 4s fills before 3d but is written after it in the shorthand. The question supplies Z up to 36, so you never need sub-shells beyond 4p.

Start from the neutral atom's electron count, which equals the atomic number, then adjust for the ionic charge. A positive ion has lost electrons, so subtract the charge from the electron count; a negative ion has gained electrons, so add the magnitude of the charge. For example, Mg (Z = 12) as Mg²⁺ has 10 electrons: 1s² 2s² 2p⁶, the same configuration as Ne. Cl (Z = 17) as Cl⁻ has 18 electrons: 1s² 2s² 2p⁶ 3s² 3p⁶, matching Ar. The scope is s- and p-block elements up to Z = 36, so you handle ions such as Na⁺, Mg²⁺, Al³⁺, O²⁻, F⁻, S²⁻ and Cl⁻. Note that removing electrons from a p-block atom takes them from the highest-energy occupied sub-shell first.