Electrons, energy levels and ionisation energy
Section: 1 Atomic Structure | Syllabus: Cambridge AS Level Chemistry (9701)
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 |
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
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:
ExamplesNa (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:
ExamplesNa: [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.
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.
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 TipSpin-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 PointThe 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 alkanes (e.g. chlorination) - studied in organic chemistry.
- The dot (•) notation is used to show the unpaired electron: Cl•, CH₃•, OH•
Ionisation Energy
First Ionisation Energy (IE₁) The energy required to remove one mole of electrons from one mole of gaseous atoms of an element, to form one mole of gaseous ions with a single positive charge. Equation for First Ionisation EnergyX(g) → X⁺(g) + e⁻ ΔH = IE₁
State symbols are essential - ionisation energy is always defined for gaseous species.
Key PointIonisation energy is always endothermic (positive ΔH) - energy must be supplied to overcome the attraction between the nucleus and the electron.
Successive Ionisation Energies
Successive ionisation energies represent removing electrons one at a time from an increasingly positive ion:
Example: AluminiumIE₁: Al(g) → Al⁺(g) + e⁻
IE₂: Al⁺(g) → Al²⁺(g) + e⁻
IE₃: Al²⁺(g) → Al³⁺(g) + e⁻
Factors Affecting Ionisation Energy
The ionisation energy of an element depends on the attraction between the nucleus and the outer electron. Four key factors influence this:
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