Ionic vs Covalent Bonding — Key Differences Explained
Question
What are the key differences between ionic and covalent bonding? Compare them on the basis of formation, properties, and examples. Use NaCl (ionic) and H₂O (covalent) to illustrate.
Solution — Step by Step
Step 1: How Each Bond Forms
Ionic bond: Forms by complete transfer of electrons from one atom to another. The atom that loses electrons becomes a positively charged cation; the atom that gains electrons becomes a negatively charged anion. The two oppositely charged ions attract each other electrostatically.
NaCl example: Na (2,8,1) → loses 1e⁻ → Na⁺ (2,8) [noble gas configuration of Ne] Cl (2,8,7) → gains 1e⁻ → Cl⁻ (2,8,8) [noble gas configuration of Ar] Na⁺ + Cl⁻ → attracted by electrostatic force → ionic bond
Covalent bond: Forms by sharing electron pairs between atoms. Neither atom completely transfers electrons — both nuclei attract the shared electrons, holding the atoms together.
H₂O example: O has 6 valence electrons — needs 2 more for an octet H has 1 valence electron — needs 1 more for a duet O shares 1 electron pair with each of 2 H atoms → 2 covalent bonds (O–H)
Step 2: When Does Each Type Form?
Ionic bond: Forms when the electronegativity difference (ΔEN) between two atoms is large (generally > 1.7). Typically a metal reacts with a non-metal.
- NaCl: EN(Na) = 0.9, EN(Cl) = 3.0, ΔEN = 2.1 → ionic
- MgO: ΔEN = 3.5 − 1.2 = 2.3 → ionic
Covalent bond: Forms when ΔEN is small (< 1.7). Typically two non-metals react.
- H₂O: EN(O) = 3.5, EN(H) = 2.1, ΔEN = 1.4 → polar covalent
- H₂: ΔEN = 0 → pure non-polar covalent
- CH₄: ΔEN = 2.5 − 2.1 = 0.4 → nearly non-polar covalent
Comprehensive Comparison Table
| Property | Ionic Compounds | Covalent Compounds |
|---|---|---|
| Bond formation | Electron transfer | Electron sharing |
| Bonding units | Oppositely charged ions | Discrete molecules (or network) |
| Electronegativity difference | Large (> 1.7) | Small (< 1.7) |
| Physical state (RT) | Solid | Solid, liquid, or gas |
| Melting point | High (600–2000°C) | Low (for molecular; very high for network) |
| Boiling point | High | Low (for molecular) |
| Electrical conductivity (solid) | None (ions fixed) | None |
| Electrical conductivity (liquid/solution) | Yes (free ions) | Mostly no |
| Solubility in water | Generally soluble | Varies (polar: soluble; non-polar: insoluble) |
| Solubility in organic solvents | Generally insoluble | Generally soluble (non-polar) |
| Bond directionality | Non-directional | Directional (specific bond angles) |
| Shape | Crystal lattice | Definite molecular geometry |
| Example | NaCl, MgO, CaCl₂ | H₂O, CH₄, CO₂, HCl |
Why This Works — The Physical Properties Explained
Why do ionic compounds have high melting points? In a crystal lattice, every ion is surrounded by multiple oppositely charged ions. You're not breaking one bond — you're disrupting an entire 3D electrostatic network (lattice energy of NaCl = −786 kJ/mol). It takes a lot of energy to pull those ions apart.
Why don't ionic solids conduct electricity? Electrical conduction requires mobile charge carriers. In solid NaCl, ions are locked into fixed positions in the crystal lattice — they cannot move. In molten NaCl or in aqueous solution, the ions are free to migrate toward electrodes → conduction occurs.
Why do molecular covalent compounds have low melting points? Covalent bonds within a molecule are strong (C–H bond: 413 kJ/mol). But the forces between molecules (van der Waals, dipole–dipole) are weak. To melt a molecular solid, you only need to overcome inter-molecular forces, not break the covalent bonds. That requires much less energy → low melting point.
📌 Note
There is an exception: network covalent solids (diamond, SiO₂, SiC) have covalently bonded atoms extending throughout the solid — no discrete molecules. These have extremely high melting points (diamond: 3500°C) because you must break actual covalent bonds to melt them. Don't confuse network covalent with molecular covalent.
Why This Works — NaCl vs H₂O Side by Side
| Property | NaCl (Ionic) | H₂O (Covalent) |
|---|---|---|
| Bonding | Na⁺ and Cl⁻ attracted electrostatically | O shares electron pairs with 2 H atoms |
| State at room temperature | Solid (crystalline) | Liquid |
| Melting point | 801°C | 0°C |
| Boiling point | 1413°C | 100°C |
| Electrical conductivity (liquid) | Yes (ions free to move) | No (no ions; pure H₂O has very low conductivity) |
| Shape | Giant ionic lattice (no discrete molecule) | Bent molecular shape, 104.5° |
| ΔEN | 2.1 (Na–Cl) | 1.4 (H–O) |
Alternative Method — Quick Identification Using Bonding Partners
In exams, quickly identify bond type from the elements involved:
- Metal + Non-metal → usually ionic (NaCl, MgO, CaF₂, Al₂O₃)
- Non-metal + Non-metal → usually covalent (H₂O, CO₂, NH₃, HCl, CH₄)
- Metalloid + Non-metal → often covalent or polar covalent (SiCl₄, B₂H₆)
This rule has exceptions (e.g., AlCl₃ is partially covalent), but for CBSE and NEET it's reliable for straightforward MCQs.
Common Mistake
⚠️ Common Mistake
Mistake: Thinking ionic compounds always dissolve in water and are always soluble, while covalent compounds are always insoluble.
Reality:
- Many ionic compounds are insoluble in water (AgCl, BaSO₄, CaCO₃)
- Polar covalent compounds are soluble in water (HCl, NH₃, glucose, ethanol)
- Non-polar covalent compounds are insoluble in water but soluble in organic solvents (CCl₄, benzene, oils)
The rule is: "like dissolves like" — polar solvents dissolve polar/ionic substances; non-polar solvents dissolve non-polar substances. Water is a polar solvent, so it dissolves ionic and polar covalent compounds, but not non-polar covalent compounds.