Electrons, energy levels and ionisation energy
Section: 1 Atomic Structure | Syllabus: Cambridge AS Level Physics 9702
Shells, Sub-shells and Orbitals Shell (Principal Energy Level) A main energy level occupied by electrons, described by the principal quantum number n (n = 1, 2, 3...). Higher n means further from the nucleus and higher energy.
Sub-shell A subdivision of a shell, labelled s, p, d or f. Each shell contains one or more sub-shells. Orbital A region of space around the nucleus that can hold a maximum of 2 electrons (with opposite spins).
Each orbital has a specific shape and energy. Ground State The lowest energy electronic configuration of an atom or ion, where electrons fill the lowest available energy levels first. Sub-shells: Orbitals and Electron Capacity Principal Quantum Number (n) Sub-shell Number of Orbitals Max Electrons in Sub-shell Max Electrons in Shell 1 1s 1 2 2 2 2s 1 2 8 2p 3 6 3 3s 1 2 18 3p 3 6 3d 5 10 4 4s 1 2 (32) 4p 3 6 Key Rule Each orbital holds a maximum of 2 electrons with opposite spins (↑↓).
An s sub-shell has 1 orbital, a p sub-shell has 3 orbitals, and a d sub-shell has 5 orbitals. Order of Increasing Energy (Aufbau Principle) Electrons fill sub-shells in order of increasing energy. The order is: 1s → 2s → 2p → 3s → 3p → 4s → 3d → 4p Important Exception The 4s sub-shell fills before the 3d because it is at a slightly lower energy level.
However, when a transition metal ion forms, 4s electrons are lost first - because once the 3d is occupied, the 3d energy falls below 4s. This filling order can be remembered using the diagonal arrow diagram (Aufbau diagram): draw sub-shells in a grid and follow diagonal arrows from top right to bottom left.
Figure 1.8: Aufbau Filling Order Diagram (Grid with sub-shells listed: 1s, 2s 2p, 3s 3p 3d, 4s 4p 4d 4f. Diagonal arrows running from top-right to bottom-left showing the filling order: 1s → 2s → 2p → 3s → 3p → 4s → 3d → 4p.
Highlight that 4s is filled before 3d.) Electronic Configurations Two conventions are used to express electronic configurations: Full notation Write all sub-shells explicitly with superscripts for electron numbers: Examples Na (Z=11): 1s² 2s² 2p⁶ 3s¹ Cl (Z=17): 1s² 2s² 2p⁶ 3s² 3p⁵ Fe (Z=26): 1s² 2s² 2p⁶ 3s² 3p⁶ 3d⁶ 4s² Shorthand (noble gas core) notation Use the preceding noble gas symbol in square brackets to replace the core configuration: Examples Na: [Ne] 3s¹ Cl: [Ne] 3s² 3p⁵ Fe: [Ar] 3d⁶ 4s² Ions For cations (positive ions): remove electrons from the highest n shell first.
For transition metals, remove 4s electrons before 3d. For anions (negative ions): add electrons to the next available orbital. Examples: Iron ions Fe²⁺: remove 2 electrons from 4s → [Ar] 3d⁶ Fe³⁺: remove 2 from 4s, 1 from 3d → [Ar] 3d⁵ Electrons in Boxes Notation Each box represents one orbital.
Electrons are shown as arrows (↑ = spin up, ↓ = spin down). Two key rules apply when filling orbitals in the same sub-shell: Hund's Rule: Every orbital in a sub-shell must be singly occupied before any orbital is doubly occupied.
Electrons in singly-occupied orbitals all have the same spin. Pauli Exclusion Principle: Each orbital can hold at most 2 electrons, which must have opposite spins. Example: Carbon (1s² 2s² 2p²) 1s: [↑↓] 2s: [↑↓] 2p: [↑][↑][ ] The two 2p electrons occupy separate orbitals with the same spin - not paired in one orbital.
Example: Nitrogen (1s² 2s² 2p³) 1s: [↑↓] 2s: [↑↓] 2p: [↑][↑][↑] Figure 1.9: Electrons in Boxes for H through Ne (Table showing electrons-in-boxes notation for elements H to Ne. Each row shows element symbol, configuration, and boxes for 1s, 2s, 2px, 2py, 2pz filled with up/down arrows.
Highlight Hund's rule in C, N and spin-pair repulsion in O.) Exam Tip Spin-pair repulsion explains why the fourth electron in a p sub-shell must pair up with an existing electron - this costs energy, which is why there is a slight dip in ionisation energy between N and O in Period 2.
Shapes of s and p Orbitals s orbitals are spherical in shape and are symmetrical about the nucleus. The 1s orbital is smaller than the 2s orbital. p orbitals are dumbbell (figure-of-eight) shaped, with two lobes on either side of the nucleus.
There are three p orbitals per sub-shell, oriented along the x, y and z axes (pₓ, pᵧ, p_z) - they are mutually perpendicular. Figure 1.10: Shapes of s and p Orbitals (Left: sphere labelled 's orbital' centred on nucleus.
Right: three dumbbell shapes labelled pₓ, pᵧ, p_z oriented along their respective axes - all centred on the same nucleus and mutually perpendicular. Show the nucleus as a dot at the centre of each.) Key Point The shape of an orbital represents the region of space where there is a high probability of finding an electron (approximately 95% probability).
Orbitals do not have sharp boundaries. Free Radicals Free Radical A species (atom, molecule or ion) that has one or more unpaired electrons. Free radicals are highly reactive. A chlorine atom, Cl•, is a free radical - it has 7 electrons in its outer shell, one of which is unpaired.
Free radicals are encountered in free radical substitution reactions of alka…
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