Atomic Structure and Electron Configuration

🟢 Lite — Quick Review (1h–1d)

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Atomic structure, bonding and periodicity underpins Paper 1 (multiple choice) and Paper 2 (structured) of UNEB UACE Chemistry. The atom is built from protons (+1) and neutrons (0) in a tiny nucleus, surrounded by electrons (−1) in orbitals labelled s, p and d. An element’s atomic number Z = protons; mass number A = protons + neutrons; isotopes share Z but differ in A. Electron configuration follows the Aufbau principle (fill lowest energy first), Hund’s rule (one electron per orbital before pairing) and the Pauli exclusion principle (max 2 electrons per orbital, opposite spins). Three bond types dominate: ionic (electron transfer between metals and non-metals), covalent (shared pairs, including dative bonds where one atom donates both electrons) and metallic (delocalised electron sea). VSEPR theory predicts molecular shape from lone-pair and bond-pair repulsion. Master the trends in first ionisation energy, electronegativity and atomic radius across Periods 2–3 and down Groups I, II and VII.

🟡 Standard — Regular Study (2d–2mo)

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Atomic Structure and Electron Configuration

The Bohr model placed electrons in fixed circular shells, but the quantum mechanical model describes them as orbitals — regions of high electron probability. The four quantum numbers (n, l, m, s) define each electron. Orbitals have distinct shapes: s (spherical), p (dumb-bell, three orientations px, py, pz) and d (five orientations). For UNEB UACE you must write configurations for Z = 1 to 36, e.g. Fe (Z = 26) = 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d⁶, or in shorthand [Ar] 4s² 3d⁶. Note the anomaly: chromium is [Ar] 4s¹ 3d⁵, not 4s² 3d⁴, due to extra stability of a half-filled d sub-shell.

Ionic, Covalent and Metallic Bonding

Ionic bonding forms when electronegativity difference > ~1.7: metals lose electrons to form cations, non-metals gain them to form anions, held in a giant lattice by electrostatic attraction. Lattice energy increases with higher charge and smaller ionic radius. Covalent bonds share one (sigma, σ), two (σ + π) or three (σ + 2π) pairs; a dative bond forms when one atom (the donor) supplies both electrons, e.g. NH₃ → BF₃ forming the adduct H₃N–BF₃. Metallic bonding involves a lattice of cations in a “sea” of delocalised valence electrons — explains conductivity and malleability.

Shapes and Polarity (VSEPR)

Pairs of electrons (bonding and lone) repel to maximise separation. Memorise these shapes because UNEB asks them every year:

Pairs around central atom	Bond pairs	Lone pairs	Shape	Bond angle
2	2	0	Linear	180°
3	3	0	Trigonal planar	120°
4	4	0	Tetrahedral	109.5°
4	3	1	Trigonal pyramidal	107°
5	5	0	Trigonal bipyramidal	90°/120°
6	6	0	Octahedral	90°

Molecules like CO₂ (linear, bond dipoles cancel) are non-polar; H₂O (bent, dipoles add) is polar.

Periodicity Across Period 2 and 3

Atomic radius decreases left → right (more nuclear pull, same shell).
First ionisation energy generally increases left → right but with dips at Group 3 (2p¹ is easier to remove than 2p²) and Group 6 (3p⁴ is easier than 3p⁵ because pairing forces electrons into the same orbital).
Electronegativity increases left → right.
Down a group (I, II, VII), atomic radius increases, ionisation energy decreases, and electronegativity decreases (for metals) or increases (F > Cl > Br > I).

Intermolecular Forces

Van der Waals (London dispersion) exist between all molecules; strength rises with electron count and surface area. Permanent dipole–dipole forces act between polar molecules. Hydrogen bonding is the strongest IMF and occurs when H is covalently bonded to F, O or N — it explains the anomalously high boiling points of H₂O, HF and NH₃.

🔴 Extended — Deep Study (3mo+)

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Hybridisation and Bond Types

Hybridisation reconciles orbital theory with observed geometry. sp³ (e.g. CH₄, NH₃, H₂O) gives tetrahedral arrangement; sp² (e.g. C₂H₄, BF₃) gives trigonal planar with a delocalised π system; sp (e.g. C₂H₂, BeCl₂) gives linear geometry. Sigma bonds form by head-on overlap along the internuclear axis; pi bonds form by sideways p-orbital overlap and are weaker. A double bond = 1σ + 1π; a triple bond = 1σ + 2π. UNEB often tests dative bonds through examples like the formation of H₃O⁺ (from H₂O + H⁺), [Cu(NH₃)₄]²⁺, or Al₂Cl₆.

Edge Cases and Common Traps

NaCl is ionic, HCl is covalent despite both being metal/non-metal combinations — bonding depends on electronegativity difference, not position in the periodic table.
Polar bonds do not always give polar molecules — symmetry matters (CCl₄ has polar bonds but is overall non-polar).
Electron affinity is not the same as electronegativity; EA is energy released when an atom gains an electron, while electronegativity is a comparative tendency in a bond.
Transition metals lose 4s electrons before 3d on ionisation: Fe → Fe²⁺ is [Ar] 3d⁶, not [Ar] 4s² 3d⁴.
Lattice energy (ΔH_LE) is endothermic in the opposite direction (always positive when breaking the lattice). Magnitude trends: LiF > NaCl > KBr — because charge density (charge/radius) drives the attraction.

Worked Example

Q: Explain why the boiling point of water (100 °C) is much higher than that of hydrogen sulphide, H₂S (–60 °C), even though H₂S has more electrons.

A: Both molecules are bent (VSEPR with 2 lone pairs on the central atom) and both are polar. However, in H₂O the hydrogen is bonded directly to the highly electronegative oxygen, producing strong hydrogen bonding between molecules. In H₂S, sulphur is not electronegative enough to form hydrogen bonds — only weaker permanent dipole–dipole and London dispersion forces operate. Although H₂S has more electrons (and thus stronger dispersion forces), hydrogen bonding in water dominates, giving H₂O its much higher boiling point.

Practice Prompts

Using s, p and d orbitals, write the electron configuration of copper (Z = 29) and explain why it does not follow the expected pattern.
Compare and contrast the lattice energies of NaF and NaI, and explain the difference in terms of ionic radius and charge density.

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Atomic Structure and Electron Configuration