Chemical Bonding: Ionic, Covalent and Metallic

🟢 Lite — Quick Review (1h–1d)

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

Ionic (Electrovalent) Bonding: Formed by complete transfer of electrons from a metal to a non-metal. The metal loses electrons (oxidation), the non-metal gains them (reduction).

• Example: $\text{Na} \rightarrow \text{Na}^+ + e^-$; $\text{Cl} + e^- \rightarrow \text{Cl}^-$
• Result: $\text{Na}^+\text{Cl}^-$ (sodium chloride, NaCl)
• Properties: High melting/boiling points, conduct electricity when molten or dissolved, crystalline structure.

Covalent Bonding: Formed by sharing of electron pairs between non-metal atoms.

• Single bond: 2 shared electrons (e.g., H–H, Cl–Cl)
• Double bond: 4 shared electrons (e.g., O=O, C=C)
• Triple bond: 6 shared electrons (e.g., N≡N)
• Properties: Generally low melting points (molecular) or high (network covalent like SiO₂), do not conduct electricity.

Metallic Bonding: Positive metal ions in a sea of delocalised electrons.

• Explains: malleability, ductility, electrical conductivity, thermal conductivity.

⚡ NECO Tip: Predicting bond type: Metal + Non-metal → ionic. Non-metal + Non-metal → covalent. Metal + Metal (or alloy) → metallic. Electronegativity difference $> 1.7$ → ionic; $0.4 - 1.7$ → polar covalent; $< 0.4$ → non-polar covalent. Use Pauling scale.

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🟡 Standard — Regular Study (2d–2mo)

Standard content for NECO Chemistry students with a few days to months.

Lewis Structures (Electron Dot Diagrams)

Show all valence electrons as dots around the atomic symbol.

Examples:

• F· (7 valence electrons)
• ·O· (6 valence electrons)
• H$\cdot$ (1 valence electron)

For covalent molecules: shared pairs shown as dashes or double dots.

• H₂O: H–O–H with 2 lone pairs on oxygen (total 8 electrons around O)
• CO₂: O=C=O (no lone pairs on C, but each O has 3 lone pairs)
• NH₃: H–N–H with one lone pair on N (pyramidal shape)

The Octet Rule: Atoms tend to gain, lose, or share electrons to achieve 8 valence electrons (except hydrogen = 2).

Exceptions to the octet:

• Boron (BH₃): only 6 electrons — electron deficient
• Phosphorus in PCl₅}: 10 electrons — expanded octet
• Sulfur in SF₆: 12 electrons — expanded octet
• Noble gases: already have full outer shells

Electronegativity and Polarity:

Electronegativity increases across a period and decreases down a group (Pauling scale).

• A polar covalent bond has unequal sharing (difference in EN 0.4–1.7)
• Dipole moment ($\mu$) measures bond polarity; vector pointing from δ+ to δ−

Shapes of Molecules (VSEPR Theory):

Electron pairs	Shape	Bond angle
2 bonding, 0 lone	Linear	180°
3 bonding, 0 lone	Trigonal planar	120°
2 bonding, 2 lone	Bent/V-shaped	~109.5°
4 bonding, 0 lone	Tetrahedral	109.5°
3 bonding, 1 lone	Trigonal pyramidal	~107°
2 bonding, 3 lone	Linear	180°

Examples:

• CO₂: linear (180°), no lone pairs on C
• H₂O: bent (~104.5°), 2 lone pairs on O
• NH₃: trigonal pyramidal (~107°), 1 lone pair on N
• CH₄: tetrahedral (109.5°), no lone pairs

Intermolecular Forces (Weak Forces Between Molecules):

1. Van der Waals (London dispersion): All molecules. Temporary dipoles. Strength increases with number of electrons.
2. Dipole–dipole: Between polar molecules. Permanent dipoles attract.
3. Hydrogen bonding: Between molecules with O–H, N–H, or F–H bonds. This is why water has unusually high boiling point (100°C instead of −80°C if only London forces).

⚡ NECO Common Mistakes:

• Confusing intramolecular bonds (within molecules) with intermolecular forces (between molecules)
• Thinking ionic compounds conduct electricity when solid — they only conduct when molten or dissolved
• Forgetting that covalent molecules like SiO₂ (quartz) have very high melting points because they form a giant covalent network
• Mixing up molecular shape vs electron pair geometry

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🔴 Extended — Deep Study (3mo+)

Comprehensive coverage for NECO and JAMB Chemistry preparation.

Born–Haber Cycle (Enthalpy of Ionic Bonding)

The enthalpy of formation of an ionic compound is the sum of several steps: $$\Delta H_f = \Delta H_{\text{atomisation}} + \Delta H_{\text{ionisation}} + \Delta H_{\text{electron affinity}} + \Delta H_{\text{lattice}}$$

For NaCl:

1. Sublimate Na(s): $\Delta H = +107$ kJ/mol
2. Ionise Na(g): $\Delta H = +496$ kJ/mol (first ionisation energy)
3. Dissociate ½Cl₂(g): $\Delta H = +122$ kJ/mol
4. Add electron to Cl(g): $\Delta H = -349$ kJ/mol (electron affinity)
5. Form lattice: $\Delta H = −784$ kJ/mol (lattice energy) $$\Delta H_f = 107 + 496 + 122 - 349 - 784 = -408 \text{ kJ/mol}$$

Covalent Bonding — Valence Bond Theory:

A covalent bond forms when atomic orbitals on two atoms overlap. The bond strength depends on the degree of overlap.

• Head-on overlap (sigma bond, σ): Stronger, allows rotation.
• Sideways overlap (pi bond, π): Weaker, prevents rotation.

Example — Ethene (C₂H₄):

• C–C bond: one σ bond (head-on $sp^2$–$sp^2$) + one π bond (sideways $p$–$p$)
• C=C is shorter and stronger than C–C single bond.

Hybridisation:

Hybridisation	Geometry	Example
$sp^3$	Tetrahedral	CH₄, NH₃, H₂O
$sp^2$	Trigonal planar	C₂H₄, BF₃
$sp$	Linear	C₂H₂, BeCl₂

Molecular Orbital Theory:

Antibonding orbitals ($\sigma^$, $\pi^$) are higher in energy than bonding orbitals ($\sigma$, $\pi$). Bond order: $$\text{Bond order} = \frac{\text{electrons in bonding} - \text{electrons in antibonding}}{2}$$

For O₂: Bond order = $(10 - 6)/2 = 2$ (double bond, paramagnetic — has 2 unpaired electrons, confirmed experimentally). For N₂: Bond order = $(10 - 4)/2 = 3$ (triple bond, diamagnetic).

Metallic Bonding — Free Electron Theory:

Metal ions are held together by attraction to a delocalised sea of electrons. This explains:

• Electrical conductivity: electrons move freely under an electric field
• Malleability: layers of ions can slide without breaking bonds
• High melting points of transition metals: d-electrons contribute to stronger metallic bonding

Alloy Types:

• Substitutional alloy: Atoms of similar size replace each other (e.g., Cu in Ag)
• Interstitial alloy: Smaller atoms occupy holes in the lattice (e.g., C in Fe → steel)

NECO/JAMB Patterns:

• NECO frequently asks: draw Lewis structures; predict molecular shape using VSEPR; identify types of bonds given electronegativity differences; explain physical properties (melting point, conductivity) based on bond type
• Be able to rank compounds by boiling point based on intermolecular forces
• Understand hydrogen bonding’s effect on water’s properties

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