Question
Compare the five major types of chemical bonds — ionic, covalent, metallic, hydrogen, and van der Waals. What determines which type of bond forms between two atoms? How does bond type affect physical properties like melting point and conductivity?
(CBSE 11 + JEE Main + NEET pattern)
Solution — Step by Step
| Bond Type | Between | Mechanism | Strength | Example |
|---|---|---|---|---|
| Ionic | Metal + non-metal | Electron transfer → cation + anion → electrostatic attraction | Strong | NaCl, MgO, CaF₂ |
| Covalent | Non-metal + non-metal | Electron sharing (equal or unequal) | Strong | H₂O, CH₄, CO₂ |
| Metallic | Metal + metal | Electron sea model — delocalised electrons | Moderate-Strong | Cu, Fe, Al |
| Hydrogen bond | H bonded to F/O/N … with lone pair on F/O/N | Electrostatic attraction between partial charges | Weak | H₂O…H₂O, DNA base pairs |
| Van der Waals | All molecules | Temporary dipole-induced dipole (London forces) | Very weak | Noble gases, I₂, CH₄ (solid) |
| Property | Ionic | Covalent | Metallic |
|---|---|---|---|
| Melting point | High (600-3000°C) | Low to moderate | Variable (Hg: -39°C to W: 3422°C) |
| Conductivity (solid) | No (ions fixed) | No | Yes (free electrons) |
| Conductivity (molten/dissolved) | Yes (ions free) | No | Yes |
| Hardness | Hard but brittle | Soft (molecular) or very hard (network) | Malleable, ductile |
| Solubility in water | Often soluble | Polar: soluble; Nonpolar: insoluble | Insoluble |
The type of bond formed depends on the electronegativity difference ():
- → predominantly ionic
- 0.4 < \Delta\text{EN} < 1.7 → polar covalent
- \Delta\text{EN} < 0.4 → nonpolar covalent
- Between identical metal atoms → metallic
This is a guideline, not an absolute rule — the actual bonding is often a blend.
graph TD
A["Chemical Bond Types"] --> B["Intramolecular - Strong"]
A --> C["Intermolecular - Weak"]
B --> B1["Ionic: e- transfer"]
B --> B2["Covalent: e- sharing"]
B --> B3["Metallic: e- sea"]
C --> C1["Hydrogen bond"]
C --> C2["Van der Waals"]
B1 --> D["NaCl: high MP, conducts when molten"]
B2 --> E["Diamond: covalent network, very hard"]
B3 --> F["Cu: conducts, malleable"]
style A fill:#fbbf24,stroke:#000,stroke-width:2px
style B fill:#86efac,stroke:#000
style C fill:#93c5fd,stroke:#000Why This Works
Bond type reflects the electronic structure of atoms. Metals have few valence electrons and low ionisation energy — they easily lose electrons (ionic bonding with non-metals) or share them in a delocalised sea (metallic bonding). Non-metals have high electronegativity — they share electrons with each other (covalent). Hydrogen bonding and van der Waals forces are secondary interactions that arise from partial charges and temporary dipoles.
The physical properties directly follow: strong ionic/covalent bonds mean high melting points, delocalised electrons mean conductivity, and weak intermolecular forces mean low boiling points.
Common Mistake
Students often classify hydrogen bonds as strong bonds on par with ionic or covalent. Hydrogen bonds are WEAK — about 10-40 kJ/mol compared to 200-400+ kJ/mol for covalent bonds. They are intermolecular forces, not true chemical bonds. However, their cumulative effect is significant — water’s high boiling point and DNA’s double helix stability both depend on numerous hydrogen bonds working together.
For exams, remember: Diamond (covalent network solid) has one of the highest melting points despite being covalent. This is because it is a network solid — every atom is covalently bonded to four others in a continuous 3D lattice. Do not confuse molecular covalent solids (like ice, low MP) with network covalent solids (like diamond, very high MP).