Atomic Structure and Bonding

🟢 Lite — Quick Review (1h–1d)

Rapid summary for last-minute revision before your WAEC exam.

Sub-atomic Particles:

Particle	Symbol	Mass	Charge	Location
Proton	p⁺	1	+1	Nucleus
Neutron	n⁰	1	0	Nucleus
Electron	e⁻	1/1836	-1	Electron cloud

Atomic Number (Z) = Number of protons Mass Number (A) = Protons + Neutrons Isotopes = Same Z, different A (e.g., ¹²C and ¹⁴C)

Electronic Configuration: Electrons fill shells starting from innermost (K=1):

• K shell: maximum 2 electrons
• L shell: maximum 8 electrons
• M shell: maximum 8 electrons (up to Ca), then 18
• N shell: maximum 32 electrons

Bohr Notation Example: For Silicon (Z=14): 2, 8, 4 or written as [Ne] 3s² 3p²

⚡ WAEC Tip: Electrons in the outermost shell are valence electrons. They determine chemical properties. Atoms with full outer shells (He: 2, Ne/Ar: 8) are chemically inert.

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🟡 Standard — Regular Study (2d–2mo)

For students who want genuine understanding.

Types of Bonding:

1. Ionic Bonding:

• Transfer of electrons from metal to non-metal
• Results in positive cations and negative anions
• Example: NaCl: Na (2,8,1) → Na⁺ (2,8) + e⁻; Cl (2,8,7) + e⁻ → Cl⁻ (2,8,8)
• Properties: High melting/boiling points, soluble in water, conduct electricity when molten/aq

2. Covalent Bonding:

• Sharing of electron pairs between non-metals
• Single bond: 2 shared electrons (e.g., H₂, Cl₂)
• Double bond: 4 shared electrons (e.g., O₂, CO₂)
• Triple bond: 6 shared electrons (e.g., N₂)
• Properties: Low melting points, often insoluble in water

3. Metallic Bonding:

• Positive metal ions in a sea of delocalised electrons
• Explains conductivity, malleability, ductility

Electronegativity and Polarity:

• Electronegativity: Tendency to attract bonding electrons
• Fluorine is most electronegative (4.0)
• If electronegativity difference > 1.7: ionic bond
• If 0.4 - 1.7: polar covalent
• If < 0.4: non-polar covalent

Dipole Moment: Polar molecules have unequal electron distribution.

• H₂O: Bent shape, polar (EN difference O-H = 1.24)
• CO₂: Linear, non-polar despite polar bonds (they cancel)
• NH₃: Trigonal pyramidal, polar

⚡ Common Mistake: Students confuse molecular polarity with bond polarity. CO₂ has polar bonds (C=O) but the molecule is linear, so the bond dipoles cancel → non-polar molecule.

Lewis Structures:

Dots represent valence electrons:

• H· + ·H → H:H (each H has 2 electrons - duet rule)
• ·N· + 3·H → H:N:H (N has 8 electrons - octet rule)
• O::O (double bond in O₂)

Exceptions to Octet Rule:

• PCl₅: Phosphorus has 10 electrons (expanded octet)
• SF₆: Sulfur has 12 electrons (expanded octet)
• BeCl₂: Beryllium has only 4 electrons (incomplete octet)
• NO, NO₂: Odd electron molecules

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🔴 Extended — Deep Study (3mo+)

Comprehensive theory for serious exam preparation.

Quantum Mechanical Model:

Four Quantum Numbers:

1. Principal (n): Shell (1, 2, 3, 4…), determines energy level
2. Azimuthal (l): Subshell (0 to n-1: s, p, d, f)
3. Magnetic (mₗ): Orbital orientation (-l to +l)
4. Spin (mₛ): Electron spin (+½ or -½)

Electron Configuration Rules:

• Aufbau Principle: Fill lowest energy orbitals first
• Hund’s Rule: For degenerate orbitals, fill each singly before pairing
• Pauli Exclusion: No two electrons can have all four quantum numbers the same

Orbital Energy Order: 1s → 2s → 2p → 3s → 3p → 4s → 3d → 4p → 5s → 4d → 5p → 6s → 4f → 5d…

Example - Iron (Z=26): 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d⁶ Or condensed: [Ar] 3d⁶ 4s²

d-Block Contraction: After filling d-orbitals, atomic radius decreases slightly. This affects properties of elements following d-block.

Hybridisation:

Type	Orbitals Mixed	Geometry	Bond Angle
sp³	1s + 3p	Tetrahedral	109.5°
sp²	1s + 2p	Trigonal planar	120°
sp	1s + 1p	Linear	180°

Bond Order:

• Bond order = (Number of bonding - Number of antibonding electrons)/2
• Higher bond order = shorter bond = stronger bond
• Bond order of 1 = single bond, 2 = double, 3 = triple

Intermolecular Forces (Van der Waals):

1. London dispersion forces: Present in ALL molecules (instantaneous dipoles)

   • Strength increases with molecular mass
   • Only force in noble gases and non-polar molecules

2. Dipole-dipole forces: Between polar molecules

   • CO₂ is linear → no dipole moment
   • H₂S is bent → has dipole moment

3. Hydrogen bonding: Special dipole-dipole when H bonded to F, O, or N

   • Explains high boiling point of H₂O, NH₃, HF
   • H₂O boils at 100°C; H₂S boils at -60°C

Solubility and “Like Dissolves Like”:

• Polar solvents dissolve ionic and polar covalent compounds
• Non-polar solvents dissolve non-polar covalent compounds
• Water is polar; oil is non-polar → oil doesn’t dissolve in water

Born-Haber Cycle (Enthalpy of Formation):

$$\Delta H_f = \Delta H_{\text{sub}} + \frac{1}{2}\Delta H_{\text{diss}} + \Delta H_{\text{IE}} + \Delta H_{\text{EA}} + \Delta H_{\text{lattice}}$$

⚡ WAEC Previous Year Pattern:

Year	Question	Concept
2023	Electronic configuration of transition metal	Aufbau principle
2022	Lewis structure of compound	Covalent bonding
2021	Type of bonding from properties	Ionic vs covalent

Shapes of Molecules (VSEPR):

Steric Number	Shape	Examples
2	Linear	CO₂, BeCl₂
3	Trigonal planar	BF₃, SO₃
4	Tetrahedral	CH₄, NH₄⁺
3	Trigonal pyramidal	NH₃, PCl₃
2	Bent	H₂O, SO₂

⚡ Exam Strategy: For bonding questions, consider physical properties first. High melting point + soluble in water + conducts electricity when molten → ionic. Low melting point + insoluble in water → covalent. Then check electronegativity difference to confirm.

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