Chemical Bonding – AIIMS Nursing, Chemistry
Why it matters for nursing
• Drugs, nutrients, and biomolecules are defined by how atoms are linked.
• Understanding bond type helps predict solubility, reactivity, and how a drug is metabolized.
• Many clinical tests (e.g., pH, ion exchange, dialysis) rely on ionic or hydrogen bonding concepts.
1. Simple Explanation
A chemical bond is the force that holds atoms together in a molecule or compound.
- Ionic bonds: Transfer of electrons → charged ions held by electrostatic attraction.
- Covalent bonds: Sharing of electron pairs → neutral or partially charged atoms.
- Metallic bonds: Delocalised “sea” of electrons shared by many metal atoms.
- Hydrogen bonds & van der Waals forces: Weaker interactions that still influence structure and function (e.g., DNA, protein folding).
2. Key Points
| Bond Type | How it forms | Typical atoms involved | Key features | Examples in biology/medicine |
|---|---|---|---|---|
| Ionic | Electron transfer | Highly electropositive (Na, Ca) + highly electronegative (Cl, F) | +/− ions, high melting/boiling points, soluble in water | Electrolytes (Na⁺, K⁺), salts used in IV solutions |
| Covalent (Non‑polar) | Share two electrons equally | Like‑halogens (Cl₂), H₂, O₂ | No net charge, low polarity | O₂, N₂, CH₄ |
| Covalent (Polar) | Share two electrons unequally | H‑O, C‑O, C‑N | Partial charges (δ⁺, δ⁻), dipole moment | H₂O, CO₂, glucose |
| Metallic | Delocalised electrons shared by lattice | Group 1‑12 metals | High conductivity, malleability, ductility | Fe, Cu, Mg (used in implants) |
| Hydrogen bond | H attached to electronegative atom (N, O, F) interacts with lone pair on another electronegative atom | H₂O, NH₃, H₂O‑protein | Weaker than covalent, directional | Stabilises DNA base pairs, protein secondary structure |
| van der Waals (London) | Temporary dipoles in all atoms | Any atoms | Very weak, short‑range | Hydrophobic interactions in cell membranes |
Bond Polarity & Electronegativity
- Electronegativity (EN): Measure of an atom’s ability to attract electrons.
- ΔEN = |EN(A) – EN(B)|
- ΔEN < 0.5 → non‑polar covalent
- ΔEN 0.5–1.7 → polar covalent
- ΔEN > 1.7 → ionic (often covalent‑ionic hybrid)
Bond Energy & Bond Order
- Bond energy: Energy required to break a bond (kJ mol⁻¹).
- Bond order: Number of bonding electron pairs (1 = single, 2 = double, 3 = triple).
- Higher bond order → stronger bond, shorter bond length.
Lewis Structures
- Represent shared valence electrons as dots or lines.
- Rules: Octet rule (except H), minimize formal charge, satisfy valence electrons.
3. Worked Example – Formation of a Water Molecule (H₂O)
| Step | Description | Diagram |
|---|---|---|
| 1. Count Valence Electrons | H: 1 each → 2; O: 6 → total 8 valence electrons. | |
| 2. Connect Atoms | Place O as central atom; form single bonds to each H (2 bonds → 2×2 = 4 electrons). | H–O–H |
| 3. Complete Octet on Central Atom | Place remaining 4 electrons as two lone pairs on O. | H–O–H with two lone pairs (••). |
| 4. Check Charges | All atoms satisfy octet; formal charges = 0. | |
| 5. Determine Polarity | O is more electronegative than H (EN_O = 3.44, EN_H = 2.20). ΔEN ≈ 1.24 → polar covalent. | |
| 6. Bond Angle | VSEPR: 2 bonds + 2 lone pairs → AX₂E₂ → bent geometry (~104.5°). | |
| 7. Resulting Properties | Partial negative on O, partial positive on H → dipole moment, high dielectric constant, good solvent for ions. |
Quick check:
- 2 H atoms → 2 bonds → 4 electrons used.
- 8 total → 8–4 = 4 left → 2 lone pairs.
- Formal charge: O: 6 valence – (6 non‑bonding + 2 bonding/2) = 0. H: 1 valence – (0 non‑bonding + 1 bonding) = 0.
4. Quick Revision Summary
| Topic | Point |
|---|---|
| Ionic bond | Electron transfer → oppositely charged ions; high solubility in water. |
| Covalent bond | Electron sharing; non‑polar vs. polar based on ΔEN. |
| Metallic bond | Del |